AMERICAN SCIENCE SERIES, BRIEFER COURSE 

AN INTRODUCTION TO THE STUDY 

JLft 
OF 

CHEMISTRY 



BT 

IRA REMSEN 

Professor of Chemistry in the Johns Hopkins University 



THIRD EDITION, REVISED AND ENLARGED 




NEW YORK 3 3 1 S' # V^ 

HENRY HOLT AND COMPANY 

1893 



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Copyright, 1886, 1893, 

BY 

Henry Holt & Co. 



ROBERT DRUMMOND, ELECTROTYPER AND PRINTER, NEW YORK. 









PREFACE TO THE FIRST EDITION. 



In preparing this book, I have endeavored to keep in 
mind the fact that it is intended for those who are begin- 
ning the study of chemistry. Instead of presenting a large 
number of facts and thus overburdening the student's 
mind, I have presented a smaller number than is usual in 
elementary courses in chemistry; but I have been careful to 
select for treatment such substances and such phenomena 
as seem to me best suited to give an insight into the na- 
ture of chemical action. Usually the mind is not allowed 
to dwell for any length of time upon any one thing and 
thus to become really acquainted with it, but is hurried on 
and is soon bewildered in the effort to comprehend what is 
presented. I cannot but believe that it is much better to 
dwell longer on a few subjects, provided these subjects are 
properly selected. 

The charge is frequently made that our elementary text- 
books on chemistry are not scientific; that is to say, that 
not enough stress is laid upon the relations which exist be- 
tween the phenomena considered, — the treatment is not sys- 
tematic. The student is taught a little about oxygen, a little 
about hydrogen, a little about nitrogen, etc.; and then a 
little about potassium, a little about calcium, etc.; and he 
is left simply to wonder whether there is any connection 
between the numerous facts offered for study. It must be 
acknowledged that there are serious difficulties in the way 
of a purely scientific treatment of chemistry, but I think 

iii 



iv PREFACE TO THE FIRST EDITION, 

that it is quite possible to treat the subject more scientifi- 
cally than is customary, and thus to make it easier of com- 
prehension to the student. I have made an effort in this 
direction in the book here offered to the public. 

In teaching chemistry, two mistakes are commonly made. 
The first is that of presenting the profoundest theories of 
the science before tho student is prepared for them. 
Hence they make little impression upon his mind, and 
he only learns to repeat words about them, without having 
any real comprehension of them. 

The other mistake is that of giving directions for ex- 
periments without making it clear to the student why 
they are performed or what they teach. The result is 
that he sees little or no connection between the subjects 
treated in the text-book and the things which he works 
with in the laboratory. 

Now, the first object of a course in science should be to 
develop a scientific habit of thought. This cannot be done 
by mere study of the theories of a science, nor by hap- 
hazard experimenting. It can only be reached by system- 
atic study of the phenomena, and by recognizing the 
connection between these phenomena and the theories. 
At the outset the best plan is to study phenomena scien- 
tifically, and afterwards speculations may be introduced to 
some extent; though, in my opinion, it is better to keep 
these decidedly subordinate in an elementary course. 

At this day it is almost superfluous to emphasize the 
great importance of laboratory work as a part of a course 
in chemistry. College authorities and school boards are 
beginning to recognize the necessity of this kind of work 
for the purpose of securing satisfactory results. A labora- 
tory can be fitted up at slight cost in which all the experi- 
ments described in this book could be performed. It is 
not necessary to wait until a complete laboratory is pro- 
vided. The accommodations needed are simple, and 
there can hardly be a college or school which could not 



PREFACE TO TEE FIRST EDITION. V 

with a little effort secure the few conveniences. Should 
there, however, be such a place, the teacher can at least 
perform the experiments described. And this he had bet- 
ter do with not more than ten or a dozen students around 
him. By constantly questioning them, and getting one or 
another to help him, or to do the work, fairly satisfactory 
results can be attained. 

If the students work in the laboratory, it is of prime 
importance that they should not be left to shift for them- 
selves. They will surely acquire bad habits of work, and 
will generally fail to understand what they are doing. A 
thorough system of questioning and cross-questioning is 
necessary in order that the work shall be successful. A 
badly-constructed piece of apparatus should not be allowed, 
and cleanliness should be insisted upon from the beginning. 
The instructor should be as watchful in the laboratory as 
in the recitation-room, and should be as exacting in regard 
to the experimental work as the teacher of languages is in 
regard to the words of a lesson. A badly-performed experi- 
ment should be considered as objectionable as a bad reci- 
tation or a badly- written exercise. When teachers of chem- 
istry acquire this feeling and work in this spirit, the edu- 
cational value of laboratory courses will be greater than it 
frequently is now. The average playing with test-tubes 
and precipitates is of questionable benefit. As it has been 
dignified by the undeserved name of scientific training, and 
put forward in place of the real thing, many thinking men 
have been led to question the value of scientific training, 
and to adhere to the old drill in grammatical forms and 
mathematical problems. I do not wonder at this, but I 
should be greatly surprised to find this state of mind con- 
tinuing after really good laboratory courses are provided. 
A slovenly laboratory course in chemistry is a poor substi- 
tute for a well-conducted course in mathematics or lan- 
guages. It behooves those who are convinced of the great 



VI PREFACE TO THE FIRST EDITION. 

advantages to be derived from good laboratory courses to 
see to it that these courses are conscientiously conducted. 

A few of the experiments described in the book cannot 
well be made by every student in the laboratory. These 
the teacher should make at all events, and he should not 
only make them, but see to it that every detail is thor- 
oughly comprehended by the student. In the directions 
for the experiments the quantities recommended are in 
some cases larger than would be desirable for each student. 
The proportions being correctly given in the book, the 
absolute quantities can be regulated by the teacher to suit 
the circumstances. 

Finally, I invite correspondence from teachers who may 
use the book, and who may experience any difficulty in its 
use. I shall gladly avail myself of any suggestion which 
may help towards making it more useful. 

The apparatus needed can be obtained from any dealer 
in chemical wares, and I have no doubt that some of the 
larger houses would furnish estimates for all that is neces- 
sary for the purpose of illustrating the course. 

I. K. 

Baltimore December 21 , 1885. 



PREFACE TO THE THIRD EDITION. 



This book has been thoroughly revised by me after seven 
years' experience with it in the laboratory and the class- 
room, and I believe that it will be found to be materially 
improved. In the work of revision I have been much 
aided by friends at home and abroad who have made valu- 
able suggestions, all of which I have endeavored to consider 
without prejudice. Special attention has been given to 
the descriptions of the experiments, with the object of 
making them as clear and as suggestive as possible. A few 
which have been found to work unsatisfactorily have been 
omitted, and a few new ones have been added, among the 
latter being some of a quantitative character which have 
proved instructive where they have been tried, and it is 
hoped that, wherever the time will permit, they may be in- 
cluded in the regular course. 

The principal changes made in the book besides those 
mentioned are : 

1. A somewhat earlier introduction of the chapter on 
the atomic theory. 

2. The presentation of the periodic law before the sys- 
tematic study of the elements is taken up, and the classifi- 
cation of the elements in accordance with this law. 

3. The addition of two chapters on some of the more 
common compounds of carbon. 

vii 



Vlll PREFACE TO TEE THIRD EDITION. 

4. The addition of a chapter on qualitative analysis. 

5. A fuller treatment of the metallic elements. 

6. The use of different type for the experiments and for 
the text; and clearer paragraphing, 

I take pleasure in thanking my assistant, Dr. W. W. 
Randall, who has contributed to the value of the book by 
helping me in reading the proofs and by preparing the 
index. 

I. R. 

Baltimore, June, 1893. 






OONtENTS, XV 

PAGE 

Acids of arsenic — Antimony — Stibine, antimoniuretted hydro- 
gen — Acids of antimony — Antimony as a base-forming element 
— Bismuth — Salts of bismuth — Bismuth nitrates — General 
remarks on the characteristics of the nitrogen group — Boron — 
Boric acid — Boric anhydride '. .... 243 



CHAPTER XVIII. 

THE CARBON GROUP : CARBON AND SILICON. 
TINANIUM — ZIRCONIUM — CERIUM— THORIUM. 

Silicon — Silicic acid — Silicon dioxide, silicic anhydride — Compar- 
ison of carbon and silicon — Rare elements of this group 258 

CHAPTER XIX. 

BASE-FORMING ELEMENTS — GENERAL CONSIDERATIONS. 

Introductory — Order in which the base-forming elements will be 
taken up — Metallic properties — Occurrence of the metals — Ex- 
traction of metals from their ores — The properties of the metals 
— Compounds of the metals — Chlorides — General properties of 
the chlorides — Oxides — Hydroxides — Decomposition of salts by 
acids and by bases — Sulphides — Qualitative analysis — Hydro- 
sulphides — Nitrates — Chlorates — Sulphates — Sulphites — Car- 
bonates — Phosphates — Silicates 261 

CHAPTER XX. 

THE POTASSIUM GROUP : 
LITHIUM, SODIUM, POTASSIUM, CESIUM, RUBIDIUM (AMMONIUM). 

G eneral — Potassium — Preparation — Properties — Compounds of 
potassium — Potassium iodide — Potassium hydroxide — Potas- 
sium nitrate — Uses of potassium nitrate — Gunpowder — Potas- 
sium chlorate — Properties — Uses — Potassium cyanide — Potas- 

* sium sulphate — Sodium— Preparation— Properties— Compounds 
of sodium — Sodium chloride — Properties — Uses — Sodium hy- 
droxide — Sodium nitrate — Sodium sulphate — Sodium thiosul- 
phate— Sodium carbonate — The Le Blanc process— The Solvay 



xvi CONTENTS. 

PAGE 

or ammonia process — Monosodium carbonate, primary sodium 
carbonate — Disodium phosphate — Sodium borate — Ammonium 
salts — Ammonium chloride— Ammonium sulphide— Ammonium 
hydrosulphide — Sodium — Ammonium phosphate — General 
characteristics of the metals of the alkalies— Rare elements of 
this group — Relations between the atomic weights of the 
members of this group — Flame reactions — The spectroscope. . . 283 

CHAPTER XXI. 

THE CALCIUM GROUP : 
CALCIUM, BARIUM, STRONTIUM, GLUCIKUM. 

General — Calcium — Compounds of calcium — Calcium chloride — 
Calcium oxide — Calcium hydroxide — Uses — Calcium hypochlo- 
rite — Properties — How bleaching-powder acts in bleaching — 
Decomposition of bleaching-powder by boiling its solution — 
Uses — Calcium carbonate — Temporary hardness — Calcium 
sulphate — Permanent hardness — Calcium phosphates — Calcium 
phosphate essential to plant-growth — Artificial fertilizers — 
Formation of calcium phosphate by precipitation — Primary 
calcium phosphate — Calcium silicate — Glass — Mortar — Cements i 
Calcium sulphide — Barium and strontium — Flame-reactions — 
Relations between the atomic weights of the members of this 
group 303 

CHAPTER XXII. 

THE MAGNESIUM GROUP : 
MAGNESIUM, ZINC, CADMIUM. 

Magnesium — Manufacture — Properties — Applications — Com- 
pounds of Magnesium — Magnesium oxide — Magnesium chlo- 
ride — Magnesi um sulphate — Uses — Zinc — Metall urgy — Proper- 
ties — Applications — Alloys — Zinc oxide — Zinc sulphate — Zinc 
chloride — Some insoluble compounds of zinc 319 

CHAPTER XXIII. 

THE COPPER GROUP : COPPER, MERCURY, SILVER. 

Copper — Metallurgy — Properties — Precipitation of copper — Ap- 
plications — Alloys — Compounds of copper — Copper forms two 
series of compounds — Cuprous oxide — Cupric oxide— Copper 



CONTENTS. xvn 

PAGE 

sulphate — Copper sulphide — Copper-plating — Mercury— Uses 
— Amalgams — Compounds of mercury — Mercuric oxide — Mer- 
curous chloride — Mercuric chloride — Mercuric sulphide — 
Precipitation of mercury as mercurous chloride — Silver — Met- 
allurgy of silver — Pattinson's method — Zinc method — Amalga- 
mation process — Properties — Alloys of silver — Compounds of 
silver — Silver nitrate — Applications of compounds of silver in 
photography — Precipitation of metallic silver — Insoluble com- 
pounds of silver — Argentous and argentic compounds — The 
specific heat of elements as a means of determining their atomic 
weights 325 



CHAPTER XXIV. 



THE ALUMINIUM GROUP : 

ALUMINIUM, GALLIUM, INDIUM, THALLIUM, SCANDIUM, YTTRIUM, 
LANTHANUM, AND YTTERBIUM. 

General — x\luminium — Preparation — Properties — Applications — 
Compounds of aluminium — Aluminium oxide — Aluminium 
hydroxide — Alums — Aluminium silicate — Natural decomposi- 
tion of feldspar — Kaolin — Clay — Ultramarine — Porcelain — 
Earthenware — Action of soluble carbonates and soluble sul- 
phides on solutions of aluminium salts — Rare elements of the 
aluminium group 339 



CHAPTER XXV. 

THE LEAD GROUP : 
LEAD, TIN, AND GERMANIUM, 

General— Lead— Metallurgy— Properties— Uses — Compounds of 
lead and oxygen— Lead oxide— Lead peroxide— Salts of lead 
—Lead acetate— Insoluble salts of lead— Lead carbonate — 
Lead sulphide —Metallurgy — Properties — Uses — Alloys — 
Stannous and stannic compounds— Stannous chloride— Stannic 
oxide— Metastannic acid -Stannic chloride— Stannic sulphide 
—How to distinguish between tin and other metals 349 



xviu CONTENTS. 



CHAPTER XXVI. 

THE IRON GROUP: 

IRON, COBALT, NICKEL. 

PAGE 

Iron— Occurrence — Metallurgy — Varieties of iron — Steel — Uses — 
Properties of iron — Iron forms two series of compounds — Fer- 
rous compounds are converted into ferric compounds by oxida- 
tion — Ferrous chloride — Ferrous sulphate — Iron alum — Fer- 
rous oxide — Ferric oxide — Ferroso- ferric oxide — Ferric acid — 
Sulphides of iron — Iron pyrites, or pyrite— Nickel — Cobalt.. . . 358 

CHAPTER XXVII. 

MANGANESE— CHROMIUM— URANIUM. 

Manganese— compounds of manganese with oxygen— Comparison 
of manganese with aluminium and with iron— Formation of 
manganous salts — Manganese dioxide — Weldon's process for 
the regeneration of manganese dioxide in the preparation of 
chlorine — Potassium permanganate — Reduction of potassium 
permanganate — Comparison of potassium permanganate with 
potassium perchlorate — Chromium — Compounds of chromium 
— Potassium chromate — Potassium dichromate — The chromate 
and dichromate are good oxidizing agents — Insoluble chromates 
— Chrome alum — Comparison of chromium with aluminium, 
iron, and sulphur — Uranium 3G9 

CHAPTER XXVIII. 

PALLADIUM— PLATINUM— GOLD. 

Palladium — Platinum — Alloys of platinum— Platinum chloride — 
Platinum bases — Gold — Forms in which gold occurs in nature 
— Metallurgy of gold — Properties — Alloys of gold — Chlorides 
of gold 376 

CHAPTER XXIX. 

SOME FAMILIAR COMPOUNDS OF CARBON. 

Organic chemistry — Occurrence of the compounds of carbon — For- 
mation of hydrocarbons — Distillation of coal — Distillation of 
wood— Distillation of bones — Fermentation — Classes of com- 
pounds of carbon— Compounds of carbon and hydrogen — Pe- 



CONTENTS. xix 



troleurn — Refining of petroleum — Hydrocarbons in petroleum — 
Homology — The ethylene series of hydrocarbons — The acety- 
lene series — The benzene series — Marsh-gas, methane, fire- 
damp — Substitution -products of the hydrocarbons — Chloroform 
— Iodoform — Ethylene, olefiant gas — Acetylene— Methyl alco- 
hol, wood -spirit — Ethyl alcohol, spirits of wine — Different 
kinds of fermentation — Distillation of fermented liquids — Prop- 
erties of alcohol — Uses of alcohol— Glycerin — Properties — 
Acetic aldehyde, ordinary aldehyde — Chloral — Formic acid — 
Acetic acid — Properties — Uses — Salts of acetic acid — Fatty 
acids — Butyric acid — Palmitic acid — Stearic acid — Soaps — Use 
of soap — Action of soap on hard waters— Relations of the soap 
industry to other industries — Oxalic acid — Lactic acid — Malic 
acid — Tartaric acid — Citric acid — Ether — Action of acids upon 
alcohols — Saponification — Fats — Butter — Ethereal salts as 
essences — Nitroglycerin — Comparison of formulas — Alcohols — 
More complex alcohols — Radicals or residues — Acids 382 

CHAPTER XXX. 

OTHER COMPOUNDS OF CARBON. 

The carbohydrates — Grape-sugar, glucose, dextrose — Formation 
of dextrose — Manufacture of dextrose or glucose — Properties — 
Levulose, fruit-sugar — Cane-sugar — Sugar-refining^Molasses 
— Properties of sugar — Sugar of milk, lactose— Bouring of 
milk— Cellulose— Properties — Gun-cotton, pyroxylin, nitrocel- 
lulose — Collodion— Celluloid — Paper — Starchy-Manufacture of 
starch — Properties — Flour -— Bread-making — Aromatic com- 
pounds—Nitrobenzene — Aniline — Aniline dyes— Phenol, car- 
bolic acid — Oil of bitter almonds, benzoic aldehyde — Benzoic 
acid — Balsams and odoriferous resins-^Gallic acid — Tannic 
acid , tannin — Tanning — Indigo— Naphthalene — Anthracene — 
Alizarin — Glucosides— Myronic acid — Alkaloids — Quinine- 
Cocaine — Nicotine — Morphine , . 402 

CHAPTER XXXI. 

QUALITATIVE ANALYSIS. 

General — Examples for practice — List of substances for examina- 
tion — Study of Group I — Study of Group II — Study of Group 
III— Study of Group IV— Study of Group V— Study of Group 
VI — General directions — Classification of substances studied. . 413 

Index , 427 



An Introduction to the Study of Chemistry. 



CHAPTER I. 



CHEMICAL ACTION.— ELEMENTS.— COMPOUNDS.— HOW TO 
STUDY CHEMISTRY. 

Introductory. — Those things which are most familiar to 
us are apt to be regarded with least wonder and to occasion 
the least thought. Take, for example, the changes in- 
cluded under the head of fire. Unless we have studied 
these changes with care, what can we make of them ? We 
see substances destroyed by fire. They apparently disap- 
pear. We feel the heat produced by the burning. We 
know that this heat disappears, and we have nothing left 
in the place of the substance burned. Take, as another 
example, the rusting of iron. W T e all know that iron when 
exposed to moist air undergoes a serious change, becoming 
covered with a reddish-brown substance which we call rust. 
If the piece of iron is comparatively thin, and it is al- 
lowed to lie in the air long enough, it will be completely 
changed to the reddish-brown substance, and no iron as 
such will be left. If a spark is brought in contact with 
gunpowder there is a flash and the powder disappears, 
dense smoke appearing in its place. What are the causes 
of these remarkable changes ? Can we learn anything 
about them by study ? 



2 INTRODUCTION TO CHEMISTRY. 

Chemical Changes. — In those changes which have been 
referred to, the substances changed disappear as such. 
After the fire, the wood or the coal, or whatever may have 
been burned, is no longer to be found. The gunpowder 
after the flash is no longer gunpowder. The rusted iron 
is no longer iron, and, no matter how long the rust may be 
allowed to lie unmolested, it will not return to the form of 
iron. Iron may, further, be changed by contact with other 
substances than air so as to lose its properties. Strong 
vinegar, which contains the substance known to chemists 
as acetic acid, acts upon iron, causing it to lose its charac- 
teristic properties. The substances known as muriatic or 
hydrochloric acid, nitric acid, and sulphuric acid also act 
upon iron and give rise to the formation of new substances 
which have not the properties of iron. Changes of this 
kind in which the substances disappear and something 
else is formed in their place are known as chemical changes, 
and chemistry is the science which has to deal with 
changes in the composition of substances. 

Physical Changes. — There are many changes taking 
place which do not affect the composition of substances. 
Iron, for example, may be changed in many ways and still 
remain iron. It may become hotter or colder. Its posi- 
tion may be changed. The difference between a piece of 
iron moving and a piece at rest is a very wonderful one, 
though we are not, as a rule, much impressed by the differ- 
ence. The iron may be struck in such a way as to cause it to 
give forth a sound. While giving forth the sound its condi- 
tion is certainly different from that in which it does not give 
forth sound. The iron may be made so hot that it gives 
light. A piece of iron may be changed further by connect- 
ing it with what is known as a galvanic battery. We can 
easily recognize the difference between a piece of iron 
through which a current of electricity is passing and one 
through which no current is passing. Finally, when a piece 



PHYSICS AND CHEMISTRY. 3 

of iron is brought in contact with a piece of loadstone, it ac- 
quires new properties. It now has the power to attract and 
hold to itself other pieces of iron. In all these cases, then, 
the iron is changed, but it remains iron. After the moving 
iron comes to rest it is exactly the same thing that it was 
before. After the iron which is giving forth sound has 
ceased to give forth sound, it returns to its original condi- 
tion. Let the heated iron alone and it cools down, ceasing 
soon to give off light, and giving no evidence of being- 
warm. Eemove the iron from contact with the galvanic 
battery and it loses those properties which are due to the 
current of electricity. In time, the iron which is mag- 
netized by contact with the loadstone loses its magnetic 
properties. Such changes are called physical changes. 

Physics and Chemistry. — From what has been said in 
regard to the kinds of change which iron can undergo, we 
see that these changes are of two kinds : 

1st. Those w r hich do not permanently affect the iron. 

2d. Those which do permanently affect the iron and 
which necessarily cause the formation of new substances 
with properties quite different from those w r hich belong to 
the iron. What is true of iron is true in general of all 
other substances. We therefore have tw r o classes of changes 
presented to us for study : 

1st. Those which do not affect the com position of sub- 
stances. 

2d. Those which affect the composition of substances 
and give rise to the formation of new substances with new 
properties. 

Changes of the first kind are called physical changes. 
Those of the second kind are called chemical changes. 

That branch of knowledge which has to deal w T ith physi- 
cal changes is known as Physics; and that which has to 
deal with chemical changes is know T n as Chemistry. 

Everything that has to do with motion, w r ith heat, light, 



4 INTRODUCTION TO CHEMISTRY. 

sound, electricity, and magnetism, is studied under the head 
of Physics. Everything that has to do with the composi- 
tion of substances and changes in the composition is studied 
under the head of Chemistry. 

Experiment 1. — Hold a piece of platinum wire in the flame of 
the laboratory burner for a moment. Remove it and hold it for 
a few moments in the air. What kind of change did it undergo 
in the flame ? Hold a piece of magnesium ribbon in the flame by 
means of a pair of pincers. What kind of change takes place ? 
Give reasons for your conclusions.— Mention some phenomena 
familiar to you that further illustrate these two kinds of change. 

Relations between the Different Kinds of Change. — Al- 
though at first sight the different kinds of change referred 
to appear to be quite distinct from one another, they are, 
in reality, closely related. If a body in motion is stopped 
suddenly, it becomes hot. Many examples of a similar 
transformation of motion into heat are familiar : a wire 
becomes hot when hammered on an anvil ; a coin rubbed 
on cloth becomes hot. In both cases the cause of the heat 
is the interference with the motion. The hammer is 
stopped and becomes hot; the coin is not stopped, but the 
motion is interfered with, and we have to push harder in 
order to move it over the cloth than we should to move it 
in the air. A wire through which a current of electricity 
is passing is heated, and if the wire is small and the cur- 
rent strong it will become so hot that it will give off light. 
Here the electricity causes heat and light. Again, we 
know that by means of heat we can produce motion. The 
steam-engine is the best example of this kind of trans- 
formation. We build a fire; this heats the water in the 
boiler; the water is converted into steam, Avhich expands 
and moves the piston; and the motion of the piston is the 
seat of all the complex motions that take place in the 
different parts of the engine. The train or the ship moves. 
What moves it ? Plainly, the heat is the cause of the 



CHEMICAL CHANGES CAUSED BY HEAT. 5 

motion. But we can go a step farther back and ask what 
causes the heat. The answer is obvious. It is the burn- 
ing of the fuel. But, in burning, the composition of the 
fuel is completely changed. A change is produced which 
in itself is not heat. When a piece of coal burns, then, it is 
undergoing a change in composition, and, as a result of 
this change, heat is produced. The heat is, therefore, pro- 
duced by a chemical change in the coal, and we may say 
that the motion of the steam-engine is the result of the 
chemical change taking place in the coal or wood which, in 
burning, produces the heat. 

Chemical Changes caused by Heat. — Just as in all 
ordinary fires we have heat produced as a result of chemi- 
cal changes in the fuel, so we may have chemical changes 
produced by heat or by electricity. 

Experiment 2.— In a clean, dry test-tube put enough white 
sugar to make a layer i to -J an inch thick. 
Hold the tube in the flame of a spirit-lamp or a 
laboratory burner, as shown in Fig. 1. — Describe 
what takes place. What is the appearance of 
the substances left in the tube ? Is there any 
sugar left ? After cooling, taste the mass. 

Experiment 3. — From a piece of hard glass 
tubing of about 6 to 7 millimetres (J inch) 
internal diameter cut off a piece about 10 cen- 
timetres (4 inches) long by making a mark 
across it with a triangular file, and then seizing 
it with both hands, one on each side of the 
mark, pulling and at the same time pressing m z ^^ 
slightly as if to break it. Clean and dry it, and 
hold one end in the flame of a laboratory burner until it melts 
together. During the melting turn the tube constantly around 
its long axis so that the heat may act uniformly upon it. Put 
into it enough red oxide of mercury (mercuric oxide) to form a 
layer about 12 millimetres (£ inch) thick. Heat the tube as in 
the last experiment. During the heating thrust into the tube a 
splinter of wood which has a spark on the end. Take it out and 




6 



INTRODUCTION TO CHEMISTRY. 



put it back again a number of times. What changes do you* 
observe in the substance in the tube ? What takes place when 
the splinter with the spark is thrust into the tube ? 

Electric Currents caused by Chemical Changes. — In 

a galvanic battery there are always substances which 
are undergoing changes in composition, and the electric 
current is due to these changes. It is therefore true that 
electric currents are produced by chemical changes. A 
simple form of a battery is represented in Fig. 2. 

The plates marked K are of copper, those marked Z of 
zinc. The plates are connected together by wires, as shown. 
In each vessel there is poured a mixture of sulphuric acid 
and water. This mixture acts upon the zinc, producing a 
chemical change in it. This is the principal cause of the 
electric current which passes through the wire. As has 




Fig. 2. 



already been stated, this wire not only conducts the electric 
current, but also becomes heated. Here, then, we have 
an electric current caused by chemical change, and heat 
caused by the electric current. 

Chemical Changes caused by the Electric Current. — This 
may be well illustrated by the action of an electric current 
on water. 

Experiment 4. — To the ends of the copper w T ires connected with 
two cells of a Bunsen's or Grove's battery fasten small platinum 
plates say 25 mm. (1 inch) long by 12 mm. (| inch) wide. Insert 
these platinum electrodes into water contained in a small shallow 



DELATION BETWEEN CHEMISTRY AND PHYSICS. 9 

Relation between Chemistry and Physics. — The experi- 
ments performed will suffice to show that the different 
kinds of changes, both the physical and the chemical, are 
more closely related to one another than they appear to be 
at first sight. In consequence of this relation we cannot 
deal with chemical changes without constantly having to 
deal with physical changes. For a thorough understanding 
of chemical changes it is necessary to have some knowledge 
of the changes produced by heat and electricity. It will 
be found that whenever chemical changes take place, heat 
changes and electric changes also take place. And it will 
be found, too, that, in order to bring about chemical 




Fig. 5. 

changes, use is frequently made of heat and electricity. 
If, therefore, the student has not studied physics, he should 
familiarize himself with a few of the elementary facts of 
the science before undertaking the study of chemistry. 
He should know what physical changes can be produced 
by heat; what boiling is; what evaporation is; what "con- 
densing a vapor '" means; what the expression "to pass 
an electric current" means; how the more common forms 
of galvanic batteries are made, etc., etc. All these matters 
are of importance in studying chemical changes, and still a 
text-book of chemistry is not the proper place to treat of 
them. It will therefore be assumed that the student has 
this knowledge. 



10 INTRODUCTION TO CHEMISTRY. 

Object of the Chemist's Study. — Everything that has to 
do with the composition of substances is the object of the 
chemist's study. Most substances can by proper methods 
be separated into simpler ones, and these again into still 
simpler ones which cannot be further decomposed by any 
means known to us. Such substances as cannot be decom- 
posed into simpler ones by us are called elements. Now, 
although there are thousands of different substances, these 
are really made up of a comparatively small number of ele- 
ments. The number of elements thus far discovered is 
between sixty and seventy, but the larger number of these 
are rare. We shall find that most things we have to deal 
with are made up of about a dozen elements, and that most 
of the chemical changes which are taking place around us, 
and which we need to study in order to get an insight into 
the nature of chemical action, take place between this small 
number of elements. 

Mechanical Mixtures. — Most of the substances found in 
nature are made up of several others. Wood, for example, 
is very complex, containing a large number of distinct sub- 
stances intimately mixed together. Some of these can be 
isolated, but it is impossible to isolate them all with the 
means at present at our command. Most rocks are also 
quite complex, and it is 'a difficult matter to isolate the 
constituents. If we look at a piece of coarse-grained gran- 
ite we see plainly enough that it contains different things 
mixed together, and if it is broken up we can pick out 
pieces of different substances from the mass. If we now 
examine a piece of each of the different substances, it ap- 
pears to be homogeneous, i.e., we cannot recognize the 
presence of more than one kind of thing in any one piece. 
If the piece is powdered, some of the powder can be ex- 
amined with a microscope without the presence of more 
than one substance being recognized. We are able to 
isolate three substances from granite by simply breaking it 



MECHANICAL MIXTURES. 11 

up. We might therefore conclude that granite consists of 
three substances. This is true, but it is not the whole 
truth. For it is possible to get simpler substances from 
each of the three. This, however, is a much more diffi- 
cult process than the separation first accomplished. Sub- 
stances must be brought in contact with each of the three 
constituents which change their composition, i.e., act 
chemically upon them, and high heat must be used to 
aid the action. It is thus possible to separate the three 
components of granite into their elements. 

Substances may then be united in different ways. They 
may be so united that it is a simple thing to separate 
them by mechanical processes. Or they may be so united 
that it is impossible to separate them by mechanical pro- 
cesses. By a mechanical process is meant any process 
which does not involve the use of heat, electricity, or 
chemical change. Thus, the mechanical process made 
use of in the case of granite consisted in picking out the 
pieces, s/ 



Experiment 7.— Mix a gram or two of powdered roll-sulphur 
and an equal weight of very line iron filings in a small mortar. 
Examine a little of the mixture with a microscope. Do you see 
both the sulphur and the iron ? 

Experiment 8. — Pass a small magnet through the mixture above 
prepared. Unless the substances used are thoroughly dry, parti- 
cles of sulphur will adhere to the magnet, but even then it will be 
seen that most of that which is taken out of the mixture is iron. 
This separation is a mechanical separation. It is a somewhat 
more refined method of picking out than that referred to in the 
case of the granite. 

Experiment 9. — Pour two or three cubic centimetres of bisul- 
phide of carbon on a little powdered roll-sulphur in a dry test- 
tube. What takes place ? Treat iron filings in the same way. 
What takes place ? Now treat a small quantity of the mixture 
with bisulphide of carbon. After shaking for some time let the 
tube stand quietly so that any solid suspended in the liquid may 



12 INTRODUCTION TO CHEMISTRY. 

settle to the bottom. Then pour off the liquid carefully on a 
watch-glass so as not to disturb the solid at the bottom. Let 
this watch-glass stand until the liquid has evaporated. Examine 
what remains undissolved in the test-tube. After the liquid has 
evaporated examine what is left on the watch-glass. What is in 
the test-tube ? What on the watch-glass ? Explain what has 
taken place. 

Experiment 10. — Make a fresh mixture of three grams each of 
powdered roll-sulphur and fine iron tilings. Grind them together 
very intimately in a dry mortar and put them in a small dry test- 
tube. Heat sharply until the mass begins to glow, then take the 
tube out of the flame. After the mass has become cool, break 
the tube, and put the contents in a mortar. Break up the solid 
and examine it. Comparing the substance with the mixture 
already examined, what differences do you note ? 

The Product is a Chemical Compound.— The new sub- 
stance formed by the action of heat on the mixture of 
sulphur and iron is no longer a mechanical mixture. AVe 
cannot decompose it by a mechanical process. The con- 
stituents are much more firmly united than they were in 
the mixture. They have lost their identity. They are 
both present, to be sure, but by means of an ordinary 
examination we cannot recognize them, as their character- 
istic properties have been lost. When the mixture began 
to glow, the act of combination began, and the glowing was 
a result of the act of combination. The sulphur and iron 
combined with each other chemically, and formed a chemical 
compound. They did not; act upon each other when simply 
brought in contact. It was necessary to heat the mixture 
in order to cause chemical combination to take place. 
The heat in this case helped the chemical action. But 
after the action began it continued without further aid 
and produced heat, as was shown by the glowing of the 
mass. 

One of the Chief Characteristics of Chemical Action. — 

The essential feature of the action in the case of iron and 



A CHARACTERISTIC OF CHEMICAL ACTION. 13 

sulphur, just studied, is this: The substances which act 
upon each other lose their individual properties and some- 
thing is formed with entirely new properties. This is true 
of every case of chemical action, and it is one of the chief 
characteristics of that kind of action. If we should exam- 
ine a number of cases of chemical action, we might be 
inclined to think that they had no common features; but 
this loss of properties and the formation of new substances 
always take place. A few examples will help to impress 
this truth. 

Experiment 11. — Examine a piece of calc-spar or of marble. 
Describe it. Heat a piece in a small glass tube, as in Experiment 
3. Does it melt ? Put a piece the size of a pea in a test-tube 
with distilled water. Thoroughly shake, and then, as heating 
usually aids solution, boil. Xow pour off a few drops of the 
liquid on a piece of platinum* foil or a watch-glass, and by 
gently heating cause the water to evaporate. If there is any- 
thing solid in solution there will be a residue on the platinum 
foil or watch-glass. If not, there will be no residue. Is the sub- 
stance soluble in water ? Now treat a small piece of the substance 
with dilute hydrochloric acid and notice what takes place. After 
the action has continued for about a minute, insert a lighted 
match in the upper part of the tube. What takes place ? Does 
the calc-spar dissolve ? To determine whether anything else has 
taken place, we shall have to get rid of the excess of hydrochloric 
acid. This we can easily do by boiling it, when it passes off in 
the form of vapor, and then whatever is in solution will remain 
behind. For this purpose put the solution in a small, clean por- 
celain evaporating-dish, and put this on a vessel containing boil- 
ing water, or a water-bath. The operation should be carried on 
in a place in which the draught is good, so that the vapors will 
not collect in the working-room. They are not poisonous, but 
they are annoying. The arrangement for evaporating is repre- 
sented in Fig. 6. 

* Platinum, an expensive metal, finds extensive use in chemical 
laboratories, for the reason that it resists the action of heat and of 
most chemical substances. 



14 



INTRODUCTION TO CHEMISTRY. 



After the liquid has evaporated and the substance in the 
evaporating-dish is dry, examine it and carefully compare its 
properties with those of the substance which was put into the 
test-tube. How does it differ in appearance from this? Is it 
harder or softer? Is it soluble in v T ater? Does it melt when 
heated in a dry tube ? Does it give off bubbles of gas when 
treated with hydrochloric acid ? Let some of it stand in contact 
with the air. What change takes place ? 



What this Experiment Shows. — The experiment shows 
that when hydrochloric acid acts upon calc-spar or marble, 

the latter at least loses its own 
properties. It might be shown 
that some of the hydrochloric 
acid also loses its properties. 
In place of the two we get a 
new substance with entirely 
different properties. The 
two substances have acted 
chemically upon each other 
and produced a chemical com- 
pound. In this case it w r as 
only necessary to bring the 
substances in contact in order 
to cause them to act chemically 
upon each other. It was riot necessary to heat them, as it 
was in the case of the iron and sulphur. /. 

Experiment 12. — Bring together in a test-tube a small piece of 
copper and some moderately dilute nitric acid. What takes 
place ? Do not inhale the gas. Describe the changes in the 
color of the liquid. Does the copper dissolve ? Examine this 
solution, as in Experiment 11, and see what has been formed. 
What are the properties of the substance found after evaporation 
of the liquid ? Is it colored ? /s it soluble in water ? Does it 
change when heated in a tube ? Is it hard or soft ? Does it in any 
w r ay suggest the copper with which you started ? 

Experiment 13. — Try the action of dilute sulphuric acid on a 




Fig. 6. 



THE CAUSE OF CHEMICAL ACTION. 15 

little zinc in a test-tube. Apply a lighted match to the mouth of 
the tube. Does the result suggest anything noticed in an experi- 
ment already performed ? What is the meaning of the bubbling 
of the liquid ? After the zinc has disappeared evaporate the 
solution as in Experiments 11 and 12. Carefully compare the 
properties of: the substance left behind with those of zinc. 

Experiment 14. — Hold the end of a piece of magnesium ribbon 
about 20 centimetres (8 inches) long in a flame until it takes fire; 
then hold the burning substance quietly over a piece of dark 
paper, so that the light white product may be collected. Compare 
the properties of this white product with those of the magnesium. 
Here again a chemical act has taken place. The magnesium has 
combined with something from the air, and heat was produced by 
the combination. The product is the white substance (compare 
Exp. 1). 

Experiment 15. — In a small, dry flask (400 to 500 ccm.) put a 
bit of granulated tin. Pour upon it 2 or 3 ccm. concentrated 
nitric acid. If no change takes place, heat gently and presently 
there will be a copious evolution of a reddish-brown gas witli a 
disagreeable smell. (Under what conditions has a gas like this 
already been obtained ?) What appears in place of the tin ? Com- 
pare the properties of the new substance with those of tin. Why 
are you justified in concluding that they are not the same thing ? 

General Conclusion. — Experiments like those just per- 
formed might be multiplied indefinitely. But a sufficient 
number have already been studied to show upon what kinds 
of observations is based the statement that : 

Whenever two or more substances act upon one another 
chemically they Jose their characteristic properties, and new 
substances with new properties are formed. 

The Cause of Chemical Action. — It is evident from what 
has already been learned that there is some power that can 
hold substances together in a very intimate way, so intimate 
that they cannot be recognized by ordinary means. AVe 
do not know what causes the sulphur and iron to combine, 
but we know that they do combine. Similarly, w r e do not 
know what causes a stone thrown upward in the air to fall 



16 INTRODUCTION TO CHEMIST BY. 

back again, but we know that it falls back. It is true, we 
may say and do say that the cause of the falling of the 
stone is the attraction of gravitation, but this does not give 
us any information, for if we ask what the attraction of 
gravitation is, we can only answer that it is that which 
causes all bodies to attract one another. We can also give 
a name to that which causes chemical combination, but 
this would not help us to understand what this cause is. 
All the chemical changes which are taking place around 
us may be referred to this cause, whatever it may be. If 
this cause should suddenly cease to operate, what would be 
the result ? Nature would be infinitely less complex than 
it now is. All substances now known to be chemical com- 
pounds would be resolved into the elements of which 
they are composed, and, as far as we know, there would be 
but about sixty or seventy different kinds of substances. 
All living things would cease to exist, and in their place 
we should have three invisible gases, and something very 
much like charcoal. Mountains would crumble to pieces. 
All water would disappear, giving two invisible gases. The 
processes of life in its many forms would be impossible, as, 
however subtle that which we call life may be, we cannot 
imagine it to exist without the existence of certain complex 
forms of matter; and, as for the life-process of larger ani- 
mals and plants, most complex chemical changes are con- 
stantly taking place within them, and these changes are 
absolutely essential to the continuation of life. 

Summary. — We have thus far learned the difference be- 
tween physical and chemical change. We have learned 
the difference between elements and chemical compounds, 
and between chemical compounds and mechanical mixtures. 
We have learned that there is a close relation between the 
different kinds of physical change and chemical change; 
and that one kind of change is capable of producing other 
kinds. We have learned how to distinguish chemical ac- 



HOW TO STUDY CHEMISTRY. 17 

tion from other kinds of action, the loss of their own 
properties which the substances suffer being a prominent 
characteristic of chemical action. 

How to Study Chemistry. — AVe might learn a great deal 
about chemical facts and learn very little in regard to the 
science of chemistry. Science is organized knowledge. As 
long as w T e do not recognize any connection between any 
set of facts observed, or as long as only a few connections 
are recognized, we cannot properly speak of a subject as a 
science. The subject must have been studied for a long- 
time. The laws governing some of the phenomena of the 
subject must have been discovered before that subject can 
be regarded, as a science. Before we can have any concep- 
tion of the science of chemistry we must become acquainted 
with some of the most important facts of the science, and 
we must also learn what connection exists between these 
facts. AVe must become familiar with substances as they 
are, but especially with the way they act upon one an- 
other. Unfortunately for our purpose, but very few simple 
substances or elements occur in the un combined form in 
nature. While, therefore, the simplest way to begin the 
study of chemical substances and chemical changes is by 
an examination of the elements, the subject is complicated 
by the fact that these elements cannot readily be obtained 
without the aid of substances which have not been studied 
and of processes which it is difficult to understand without 
some knowledge of chemistry. There are, however, two 
elements that occur in nature in enormous quantities 
and that can be obtained in the uncombined condition 
very easily. As the kinds of action which these exhibit are 
of great importance and give an insight into the nature 
of chemical action in general, the study of chemical phe- 
nomena may be profitably begun by the study of these 
two elements. They are oxygen and hydrogen. In learn- 
ing the main facts in regard to these elements much will 



18 



INTRODUCTION TO CHEMISTRY. 



be learned that will be of service in making other chemical 
phenomena comprehensible. 

The Elements and their Symbols.— Before beginning this 
study a list of the elementary substances thus far discov- 
ered is here- given. The names of those which are most 
widely distributed, and which form by far the largest part 
of the earth, are printed in small capitals. The names of 
those which are very rare are printed in italics. As has 
been stated, not more than a dozen elements enter largely 
into the composition of the earth. It has been calculated 
that the solid crust of the earth is made up approximately 
as represented in the subjoined table : 



Oxygen 44 -48.7$ 

Silicon 22.8-36.2$ 

Aluminium 9.9-6.1$ 

Iron.. . .. . t 9.9- 2.4$ 



Calcium... 6.6-0.9$ 

Magnesium 2.7-0.1$ 

Sodium 2.4-2.5$ 

Potassium . . , 1.7-3.1$ 



While oxygen forms a large proportion of the solid crust 
of the earth, it forms a still larger proportion (eight ninths) 
of water by weight, and about one fifth of the air by vol- 
ume. Carbon is the principal element entering into the 
structure of living things, while hydrogen, oxygen, and ni- 
trogen are also essential constituents of animals and plants. 
Nitrogen forms about four fifths of the air by volume. 

In representing the results of chemical action it is con- 
venient to use abbreviations for the names of elements and 
compounds. Thus, instead of oxygen we may write simply 
0, for hydrogen H, for nitrogen N", etc., etc. These sym- 
bols are used in representing what takes place when sub- 
stances act upon one another, as will be shown more clearly 
hereafter. Very frequently the first letter of the name of 
the element is used as the symbol. If the names of two 
or more elements begin with the same letter, this letter is 
used, but some other letter of the name is added. Thus, B 
is the symbol of boron, Ba of barium, Bi of bismuth, etc. 



THE ELEMENTS AND THEIR SYMBOLS. 



19 



In some cases the symbols are derived from the Latin names 
of the elements. Thus, the symbol for iron is Fe, from 
Latin ferrum; for copper, Cu, from cuprum; for mercury, 
Hg, from hydrargyrum, etc. The symbols of the more 
common elements will soon become familiar by use. It is 
not desirable to attempt to commit them to memory at 
this stage. 

The names themselves are derived from a variety of 
circumstances. Chlorine is derived from £A&po£, which 
means yellowish green, as this is the color of chlorine. 
Bromine comes from ftpcojuoz, a stench, a prominent 
characteristic of bromine being its bad odor. Hydrogen 
comes from vSoop, water, and yeveiv, to produce, signi- 
fying that it is a constituent of water. Similarly nitrogen 
comes from rirpov, nitre, and yeveiv, to produce, nitro- 
gen being one of the constituents of nitre. Potassium is 
an element found in potash, and sodium is found in soda. 



List of the Elements and their Symbols. 



Aluminium Al 

Antimony Sb 

Arsenic As 

Barium Ba 

Bismuth Bi j 

Boron B 

Bromine Br 

Cadmium Cd 

Ccesium Cs 

Calcium Ca 

Carbon C 

Cerium Ce 

Chlorine ...CI 

Chromium Cr 

Cobalt Co 

Columbium Cb 

Copper Cu 

Dtdymium Di 

Erbium E 

Fluorine F 

Gallium Ga 

Germanium Ge 

Glucinum Gl 



Gold Au 

Hydrogen H 

Indium In 

Iodine I 

Iridium ■ .Ir 

Iron Fe 

Lanthanum La 

Lead Pb 

Lithium Li 

Magnesium Mg 

Manganese Mn 

Mercury Hg 

Molybdenum Mo 

Nickel Ni 

Nitrogen X 

Osmium Os 

Oxygen O 

Palladium Pd 

Phosphorus P 

Platinum Pt 

Potassium K 

Rhodium Rh 



i 



Rubidium Rb 

Ruthenium Ru 

Scandium Sc 

Selenium Se 

Silicon Si 

Silver Ag 

Sodium Na 

Strontium Sr 

Sulphur S 

Tantalum Ta 

Tellurium Te 

Thallium Tl 

Thorium Th 

Tin Sn 

Titanium Ti 

Tungsten W 

Uranium U 

Vanadium V 

Ytterbium Yt. 

Yttrium Y 

Zinc Zn 

Zirconium Zr 



v, 



CHAPTEE II. 
A STUDY OF THE ELEMENT OXYGEN. 

In Experiment 4 it was shown that when an electric 
current is passed through water two gases are liberated. 
One of these was distinguished by the readiness with which 
substances burned in it. This gas is oxygen. A gas with 
similar properties was also obtained by heating the red ox- 
ide of mercury. This is, in fact, the same substance. 

Occurrence of Oxygen. — Oxygen is the most widely dis- 
tributed element, and it occurs also in very large quantity. 
It has been stated that it forms between forty and fifty 
per cent of the solid crust of the earth, eight ninths of 
water, and about one fifth of the air. 

Preparation of Oxygen. — The simplest way to make oxy- 
gen is by heating some substance which contains it. The 
simplest example of this kind is that furnished by the oxide 
of mercury, which when heated yields mercury and oxygen. 
If the oxide is weighed, and, after decomposition, the oxy- 
gen and the mercury are weighed, the weight of the mercury 
plus the weight of the oxygen will be found to be equal to 
the weight of the oxide. Therefore the oxide contains 
only mercury and oxygen. They are chemically combined. 
AVhen the temperature is raised sufficiently the compound 
is resolved into its elements. The chemical compound 
which contains mercury and oxygen is represented by writ- 
ing the symbols of the two elements side by side, thus, 
HgO, which signifies primarily that the two elements are 

20 



PREPARATION OF OXYGEN. 21 

in chemical combination. To represent what takes place 
when the oxide is heated this equation is used: 

HgO = Hg+0; 

which is read, mercuric oxide gives mercury and oxygen. 

Preparation of Oxygen from Potassium Chlorate.— Another 
substance which readily gives up oxygen when heated is 
potassium chlorate. This is a white, crystallized substance 
which is manufactured in large quantity and can be bought 
cheaply. It contains the elements chlorine, oxygen, and 
potassium. When heated to a sufficiently high temperature 
it gives off all its oxygen, a compound of potassium and 
chlorine being left behind. The chemical changes brought 
about in potassium chlorate by heating it are interesting, 
and they will be studied somewhat in detail a little later. 
At present they arc of interest mainly because they furnish 
the element oxygen. 

Experiment 16. — Arrange an apparatus as shown in Fig. 7. A 
represents a flask of 100 ccm. capacity. By means of a good- 
fitting cork stopper one end of the bent glass tube B is connected 
with it, and the other end, which should turn slightly upward, is 
placed under the surface of the water in C. In A put 2 to 3 grams 
(about a sixteenth of an ounce) potassium chlorate, and gently 
heat by means of the lamp. Notice carefully what takes place. 
At first the potassium chlorate will melt, forming a clear liquid. 
If the heat is increased, the liquid will appear to boil, and it will 
soon be seen that a gas is being given off. Now bring the inverted 
cylinder D filled with water over the end of the tube, and let the 
bubbles of gas rise in the cylinder. After a considerable quantity 
of gas has been collected in this way the action stops, the mass in 
the flask becomes solid, and apparently the end of the process is 
reached. But if the heat is raised still higher, gas will again 
come off, and in this second stage a larger quantity will be col- 
lected than in the first. Finally, however, the end is reached, 
and the substance left in the flask remains unchanged, no matter 
how long heat may be applied. Examine the gas as in Experi- 



22 



INTRODUCTION TO CHEMISTRY. 



merits 3 and 4. It will be shown later that in the first stage 
of the decomposition of potassium chlorate the products are 
potassium perchlorate and oxygen, and that in the second stage 
the potassium perchlorate is decomposed into potassium chloride 




Fig. 7. 

and oxygen, so that the final products of the action are potas- 
sium chloride and oxygen. 

Oxygen from Manganese Dioxide. — Another good method 
of preparing oxygen consists in heating Jblack oxide of 
manganese. This is a compound found in nature, called 
by mineralogists pyrolusite, and by chemists manganese 
dioxide. It consists of the elements manganese and 
oxygen. When this substance is heated it loses part of 
its oxygen, and there is left behind another compound 
of manganese and oxygen containing the elements in 
different proportions. 

Experiment 17. — Make some oxygen by heating to redness 4 to 
5 grams (about an eighth of an ounce) of manganese dioxide in 
a hard-glass tube closed at one end and connected at the other 
end by means of a cork with a bent glass tube. 

Oxygen from Potassium Chlorate and Manganese Dioxide. 

— The most convenient way to make oxygen in the labora- 



PREPARATION OF OXYGEN. 



23 



tory is to heat a mixture of equal parts by weight of potas- 
sium chlorate and manganese dioxide. This mixture gives 
off oxygen very readily with the aid of gentle heat. The 
potassium chlorate gives up its oxygen under these circum- 
stances. The manganese dioxide takes part in the decom- 
position, but remains behind finally in its original form. 
The chemical changes involved are quite complicated and 
cannot be studied profitably at this stage. 

Experiment 18. — Mix 25 to BO grams (or about an ounce) of 
coarsely powdered potassium chlorate with an equal weight of 




Fig. 8. 

coarsely powdered manganese dioxide in a mortar. The sub- 
stances should not be in the form of powder. Test the mixture 
by heating a very small quantity of it in a dry test-tube. If the 
decomposition takes place quietly, put the mixture in a flask, 
arranged as shown in Fig. 7, heat it, and collect the gas by dis- 



24 INTRODUCTION TO CHEMISTRY. 

placement of wajter in appropriate vessels, — cylinders, bell glasses, 
bottles with wide mouths, etc. It will also be well to collect some 
in a gasometer, such as is commonly found in chemical labora- 
tories, the essential features of which are represented in Fig. 8. 
It is made either of metal or of glass. The opening at d can be 
closed by means of a screw-cap. In order to fill it with water 
open the stop-cocks and pour the water into the upper part of the 
vessel after having screwed the cap on to d. When it is full, 
water will flow out of the small tube e. Now close all the stop- 
cocks, and take the cap from d. The water will stay in the vessel 
for the same reason that it will stay in the cylinder inverted with 
its mouth below water. To fill the gasometer with gas, put it 
over a tub or sink and introduce the tube from which gas is 
issuing into the opening at d. The gas will rise and displace the 
water, which will flow out at d. When full, screw the cap on. 
We have now a supply of gas which we can draw upon as we may 
need it. To get the gas out of the gasometer, attach a rubber 
tube to e, pour water into the upper part of the gasometer, open 
the stop-cock a and that at e, when the gas will flow out, and the 
current can be regulated by means of the stop-cock at e. 

Physical Properties of Oxygen. — In the first place, the 
gas is invisible. The slight cloud which appears in the 
vessels when the gas is first collected is due to the presence 
of a very small quantity of a substance which is not oxygen. 
If the vessels are allowed to stand for a few minutes the 
cloud will disappear, and the vessels will look the same as 
if they were filled with air. The gas is tasteless and in- 
odorous. 

Experiment 19.— Inhale a little of the gas from one of the 
small bottles. 

Oxygen is slightly heavier than the air. This can be de- 
termined by weighing a globe filled with air, then driving 
out the air by passing a current of oxygen through it for 
some time, and weighing it again. If these weighings are 
carefully made, it will be found that the relation between 
the weights of equal volumes of air and oxygen is 1 : 1.1056. 
Or, in other words, if a certain volume of air weighs 1 



PROPERTIES AND CONDUCT OF OXYGEN. 25 

gram, the same volume of oxygen will weigh 1.1056 grams. 
When oxygen is subjected to very strong pressure and a 
very low temperature, it becomes liquid. 

The properties of oxygen to which reference has thus far 
been made are its physical properties. These are its ap- 
pearance, taste, smell, relative weight, and changes in its 
condition, which still leave it in the elementary form im- 
combined chemically. Our knowledge of oxygen must, of 
course, include a knowledge of its physical properties, but, 
from the chemical point of view, it is more important for 
us to know how oxygen acts chemically. What chemical 
changes is it capable of bringing about ? What conditions 
are necessary in order that it may act chemically ? What 
laws govern the action ? What products are formed ? 

Chemical Conduct of Oxygen. — In order to get an idea of 
the way in which oxygen acts upon some simple substances 
under ordinary circumstances, we may perform a few ex- 
periments. 

Experiment 20.— Turn three of the bottles containing oxygen 
with the mouth upward, leaving them covered with glass plates. 
Into one introduce some sulphur in a so-called deflagrating-spoon, 
which is a small cup of iron or brass attached to a stout wire 
which passes through a round metal plate, usually of tin. (See 
Fig. 9.) In another put a little charcoal (carbon), and in a third 
a piece of phosphorus* about the size of a pea. Let them stand 
quietly and notice what changes, if any, take place. Sulphur, 
carbon, and phosphorus are elements, and oxygen is an element. 
It will be noticed that the sulphur and the carbon remain un- 
changed, while some change is taking place in the vessel contain- 
ing the phosphorus, as is shown by the appearance of white fumes. 

* Phosphorus should be handled with great care. It is always 
kept under water, usually in the form of sticks. If a\ small piece* is 
wanted, take out a stick with a pair of forceps, and put it under 
water in an evaporating-dish. While it is under the water, cut off a 
piece of the size wanted. Take this out by means of a pair of forceps, 
lay it for a moment on a piece of filter- paper, which will absorb most 
of the water, then quickly put it in the spoom 



23 



INTRODUCTION TO CHEMISTRY. 



After some time the phosphorus will disappear entirely, the fumes 
will also disappear, and there will be nothing to show us what has 
become of the phosphorus. If the temperature of the room is 
rather high, it may happen that the phosphorus takes fire. If it 
should, it will burn with an intensely bright light. After the 
burning has stopped, the vessel will be filled with white fumes, 
but these will quickly disappear, and the vessel will apparently be 
empty. 

What these Experiments Show. — These experiments 
show us that oxygen does not act upon sulphur and carbon 
when brought in contact with them at the ordinary tem- 
perature, and that the action upon phosphorus is generally 
slight. We might perform experiments of this kind with 
a great many substances, and we should reach the con- 
clusion that at ordinary temperatures oxygen does not act 
upon most substances. 



Action of Oxygen at Higher Temperatures. — If, however, 
the substances are heated before they are introduced into 
the oxygen, the results will be entirely different. Instead 
of conducting itself as an inactive element, oxygen will act 
with great ease upon many substances. Things such as 
coal, wood, etc., which we know wall burn in the air, burn 
in oxygen much more readily, and several substances such 

as iron, copper, etc., which will 
not burn in the air, burn in oxy- 
gen w r ith ease. 

Experiment 21. — In a cleflagrating- 
spoon set fire to a little sulphur and 
let it burn in the air. Notice whether 
it burns with ease or with difficulty. 
Notice the odor of the fumes which 
are given off. Now set fire to another 
small portion and introduce it in a 
spoon into one of the vessels containing oxygen, as shown in 
Fig. 9. It will be seen that the sulphur burns much more readily 




Fig. 9. 



ACTION OF OXYGEN AT HIGHER TEMPERATURES. 27 

in the oxygen than in the air. Notice the odor of the fumes 
given off. Is it the same as that noticed when the burning 
takes place in the air ? 

Experiment 22. — Perform similar experiments with charcoal. 

Experiment 23.— Burn a piece of phosphorus not larger than 
a small pea in the air and in oxygen. In the latter case the light 
emitted from the burning phosphorus is so intense that it is pain- 
ful to some eyes to look at it. It is better to be cautious. The 
phenomenon is an extremely brilliant one. The walls of the 
vessel in which the burning takes place become covered with a 
white substance which afterwards gradually disappears. 

Experiment 24. — Straighten a steel watch-spring* and fasten 
it in a piece of metal, such as is used for fixing a deflagrating- 
spoon in an upright position ; wind a little thread around the 
lower end, and dip it in melted sulphur. Set fire to this and in- 
sert it into a vessel containing oxygen. For a moment the sul- 
phur will burn as in Experiment 21; but soon the steel begins to 
burn brilliantly, and the burning continues as long as there is 
oxygen left in the vessel. Xotice that in this case there is no 
flame, but instead very hot particles are given off from the burn- 
ing iron. The phenomenon is of great beauty, especially if ob- 
served in a dark room. The Avails of the vessel become covered 
with a dark reddish brown substance, some of which will also be 
found at the bottom in larger pieces. This substance is a com- 
pound of iron and oxygen known as magnetic oxide of iron. 

What has Taken Place?— What lias taken place in these 
experiments? In the first place, the substances were 
simply heated before being introduced into the oxygen. 
Nothing was added to them except heat. It is clear that 
while oxygen does not act upon these substances at ordi- 
nary temperatures, it does act upon them at higher tem- 
peratures. But what does the action consist in ? We can 
determine this only by a careful study of the substances 

* Old watch-springs can generally be had of any watch maker or 
mender for the asking. A spring can be straightened by unrolling 
it, attaching a weight, and suspending the weight by the spring. 
The spring is then heated to redness from one end to the other by a 
Bunsen burner. 



28 INTRODUCTION TO CHEMISTRY. 

before and after the action. We must know not only what 
substances are brought together, but also ivliat weight of 
each; and we must know what substances are left behind, 
and the exact weights of these. In the cases mentioned it 
would be a difficult matter for one not very thoroughly 
trained in the use of chemical methods to make all these 
determinations accurately, and unless they were made accu- 
rately they would fail to furnish the desired explanation. 
The determinations have fortunately been made so fre- 
quently that there can be no doubt as to what would be 
found were the experiments to be repeated, and for the 
present it will be necessary to accept the results, and use 
them as the basis of our reasoning. Something, however, 
may be learned with but little difficulty. If in the experi- 
ment with sulphur the spoon is examined after the burn- 
ing stops, it will be found that the sulphur has disappeared. 
It will also be noticed that there is present an invisible* 
substance which has a strong, disagreeable odor. This 
substance is not oxygen and it is not sulphur, but it is a 
gas which is formed by the burning of sulphur in oxygen. 
What has become of the oxygen ? That it is no longer 
present in its original condition may be shown by intro- 
ducing a burning stick into the vessel. Instead of con- 
tinuing to burn with increased activity, as we have seen it 
do in oxygen, it is extinguished. 

In the experiment with carbon the results are similar, 
only the invisible substance has no odor. 

In the experiment with phosphorus the white substance 
which is deposited on the walls of the vessel is not phos- 
phorus, as is clear from the fact that it dissolves in water. 

Proof that the Oxygen Combines with the Burning Sub- 
stance. — The oxygen being invisible, it is more difficult to 
determine whether it enters into combination or not, but 

* The fumes first noticed subside if a little water is in the bottle, 



THE PART OXYGEN PLATS IN BURNING. 



29 



that it does can be shown by properly devised experiments. 
It is only necessary to burn a substance in a closed vessel 
containing oxygen, and to determine, after the burning, 
whether there is less oxygen than there was before. La- 
voisier, who first showed what part the oxygen plays in 
burning, made a very important experiment much like the 




Fig. 10. 

following: Some phosphorus is enclosed in a sealed tube 
with oxygen, and, by heating from without, the phosphorus 
is set on fire. After the action is over, one end of the tube 
is broken off under water, when water rushes in, showing 
that the gas that was in the tube has disappeared. A 
modification of this experiment is here described. 



80 INTRODUCTION TO CHEMISTRY. 

Experiment 25. — Arrange an apparatus as shown in Fig. 10. 
A is a glass tube about 60 cm. (2 feet) long and about 3| cm. (1| 
inches) in diameter. This is connected by means of a bent tube 
with the small flask B, of 50 to 100 ccm. capacity, which is fitted 
with a stopper having two holes. This flask is carefully dried, and 
then a thin layer of iron-dust or fine iron -filings is put on the 
bottom. The lower end of A is immersed to the extent of about 
5 cm. (2 inches) in water. A current of oxygen is now passed 
through the apparatus by connecting at C with a generator or 
gasometer. When the air has thus been displaced the current of 
oxygen is stopped, and the stop-cock at the end of C is closed. 
Now heat the iron gently by applying a flame to the flask. When the 
iron begins to glow, remove the flame. What evidence is furnished 
that the oxygen enters into combination and disappears as a gas ? 

Products Formed. — Experiments of the kind described 
have shown that, wdienever a substance burns in oxygen, 
both the substance and the oxygen lose their characteristic 
properties, and that something else is formed in their 
place. In other words, the process is one of chemical 
combination. Sulphur combines with oxygen to form a 
gaseous product known as sulphur dioxide. It is this gas 
that gives the strong odor when sulphur burns. Carbon 
combines with oxygen to form the invisible gas carbon 
dioxide, commonly called carbonic acid gas. Phosphorus 
combines with oxygen to form a white solid, phosphorus 
pentoxide, which gradually dissolves in the water present 
and disappears. All these products are well known, and 
they will be studied when the elements sulphur, carbon, 
and phosphorus are taken up. 

Proportions by Weight in which the Substances Combine 
with Oxygen. — The next question that naturally presents 
itself is this: In what proportions by weight do the sub- 
stances combine with oxygen ? Is there anything definite 
in these proportions, or do they combine in any possible 
proportions ? This is a very important question, and it has 
given rise to a great deal of experimenting, especially in 



WEIGHTS OF THE COMBINING SUBSTANCES. 31 

the early part of this century. It is impossible to repeat 
these experiments here, but the method of work can be 
made clear by a general account. Suppose magnesium is 
taken for experiment. A small quantity is accurately 
weighed by a chemical balance. It is now heated in 
oxygen, and, after the action is complete, the product is 
weighed. The experiment is repeated a number of times, 
and all the weights are carefully recorded. If every precau- 
tion is taken to secure accuracy, it will be found that 
these elements always combine in the same proportion by 
weight: 1 gram of oxygen combines with 1+ grams of 
magnesium. By similar experiments, it has been shown 
that whenever carbon burns in oxygen these two elements 
combine in the same proportion by weight — 1 gram of 
carbon combining with 2| grams of oxygen; and similar 
results have been obtained in all other cases. This act of 
combining with oxygen is one involving the action of defi- 
nite weights of substances. 

Relation of the Weight of the Product to the Weights of 
the Combining Substances. — In the experiment with mag- 
nesium described in the preceding paragraph the oxygen 
was not weighed. The increase in the weight of the 
magnesium caused by combination with oxygen w» de- 
termined, and the increase was ascribed to oxygen. A 
thoroughly satisfactory experiment would, however, in- 
volve the weighing of the magnesium, of the oxygen, and 
of the product formed. Such experiments have been 
made in great number, and it has been shown that the 
weight of the substance burned plus that of the oxygen 
used up is exactly equal to the weight of the substance 
formed. 

Burning in the Air. — One cannot well help noticing a 
strong resemblance between the burning of substances in 
oxygen and in the air, and the question naturally suggests 



32 IKTRODtJCTION TO CBEMISTk?, 

itself, Are these two acts the same ? The only way to &n* 
swer this question is to burn the same things in pure oXygeii 
and in air, and to see whether the same product Is formed 
in each case, and at the same time whether anything else is 
formed. If this comparison should be made in any case it 
would be found that whether a substance burns in the air or 
in pure oxygen the same product is formed, and nothing else. 
It is therefore certain that the act of burning in the air is 
due to the presence of oxygen. We shall learn later that the 
reason why substances do not burn as readily in the air as 
in pure oxygen is that in the air there is present a large 
quantity of another gas which does not act upon the sub- 
stances at all. 

Combustion. — By the term combustion in its broadest 
sense is meant any chemical act which is accompanied by 
an evolution of light and heat. Ordinarily, however, it is 
restricted to the union of substances with oxygen as this 
union takes place in the air, with evolution of light and 
heat. Substances which have the power to unite with oxy- 
gen are said to be combustible, and substances which have 
not this power are said to be incombustible. Most elements 
combine with oxygen under proper conditions, and are 
thei^fore combustible. Most compounds formed by the 
union of oxygen with combustible substances are incom- 
bustible. For example, the sulphur dioxide, carbon di- 
oxide, and phosphorus pentoxide obtained in Experiments 
21, 22, and 23 are incombustible. They contain oxygen, 
and they cannot directly combine with any more. 

Kindling Temperature. — We have seen that substances 
do not usually combine with oxygen at ordinary tempera- 
tures, but that in order to effect the union the temperature 
must be raised. If this were not the case, it is plain that 
every combustible substance in nature would burn up, 
for the air supplies a sufficient quantity of oxygen for this 



BLOW OXIDATION. 33 

purpose. Some substances need to be heated to a high 
temperature before they will combine with oxygen; others 
require but very slight elevation. If we were to subject a 
piece of phosphorus, of sulphur, and of carbon to the 
same gradual rise in temperature, we should find that the 
phosphorus takes fire very easily, only a slight elevation of 
temperature being necessary ; next in order would come 
the sulphur ; and last the carbon. If we were to repeat 
these experiments a number of times, we should find that 
the phosphorus would always take fire at the same tem- 
perature, and a similar result would be reached in the cases 
of sulphur and carbon. Every combustible substance 
has its kindling temperature ; that is, the temperature at 
which it will unite with oxygen. Below this temperature it 
will not unite with oxygen. If a piece of wood could be 
heated to its kindling temperature all at once, it would 
burn up as rapidly as it could secure the necessary oxygen; 
but the burning does not usually take place rapidly, for 
the reason that only a small part of it is at any one time 
heated to the kindling temperature. Watch a stick of wood 
burning, and see how, as we say, " the fire creeps " slowly 
along it. The reason of the slow advance is simply this : 
only those parts of the stick that are nearest the burning 
part become heated to the kindling temperature. 

Slow Oxidation. — Substances may combine slowly with 
oxygen without evolution of light. Thus, if a piece of iron 
is allowed to lie in moist air, it becomes covered with rust. 
This rust is similar to the substance formed when iron is 
burned in oxygen. Both are formed by the union of iron 
and oxygen. Magnesium burns in the air, as we have seen, 
and forms a white compound containing oxygen. It burns 
with increased brilliancy in oxygen, forming the same com- 
pound. If left in moist air for some days or weeks, it be- 
comes covered with a layer of the same white substance. If 
this is scraped off and the magnesium be further exposed, 



84 INTRODUCTION TO CHEMISTRY. 

it will again become covered with a layer of the compound 
with oxygen, and this may be continued until the magne- 
sium has been completely converted into the same substance 
that is formed when it burns in oxygen or in the air. Many 
other similar cases of slow oxidation might be described, 
some of which, such as the decay of wood, are constantly 
taking place in nature. 

Breathing. — The most important illustration of slow oxi- 
dation is that which takes place in our bodies, for, as we 
shall see, the food which we partake of undergoes a great 
many changes; some of the substances uniting with oxygen, 
and thus keeping up the temperature of our bodies. This, 
however, is done without evolution of light and without 
marked evolution of heat. We take large quantities of 
oxygen into our lungs in breathing. This acts upon vari- 
ous substances presented to it, oxidizing them to other 
forms which can easily be got rid of. More will be said in 
regard to the breathing process of animals and plants when 
the subject of carbon and its compounds with oxygen is 
considered. 

Heat of Combustion. — What is the chief difference be- 
tween combustion, as we ordinarily understand it, and slow 
oxidation ? So far as we can judge by a cursory examina- 
tion, it is that in the former light and heat are produced, 
while in the latter no light and very little or no heat is 
produced. Remembering that the reason why a body gives 
light is that it is heated to a sufficiently high temperature, 
the problem resolves itself into a question of heat. What dif- 
ference, if any, is there between the quantity of heat given 
off when a substance burns and when it undergoes slow 
oxidation without evolution of light ? The answer is of the 
highest importance. There is no difference whatever. In 
the one case the heat is all given off in a short space of time, 
and therefore the temperature of the substance becomes 



1 

HOW THE QUANTITY OF HEAT IS MEASURED. 35 

high and it emits light. In the other the heat is evolved 
slowly and continues for a much longer time, and therefore 
the temperature of the substance does not get very high, as 
surrounding substances conduct off the heat as rapidly as 
it is evolved. If, however, we were to measure the quantity 
of the heat, we should find it to be the same in both cases. 

How the Quantity of Heat is Measured. — The quantity of 
heat given off in a chemical reaction can be measured by 
allowing the reaction to take place in a vessel called a 
calorimeter, so constructed as to prevent loss of heat, and 
containing a known weight of water. The temperature 
of the water is noted at the beginning of the operation and 
at the end. A quantity of heat is generally stated by giv- 
ing the number of grams of water which it will raise one 
degree (Centigrade) in temperature. The quantity of heat 
necessary to raise a gram of water one degree (Centigrade) 
in temperature is the unit used in heat-measurement. It 
is called the calorie. If, therefore, we say the quantity of 
heat evolved in any reaction is 250 calories (written gener- 
ally 250 cal.), we mean simply a quantity of heat which 
would raise the temperature of 250 grams of water one de- 
gree (Centigrade) in temperature. 

To repeat, then : by the heat of combustion of a substance 
we mean simply the quantity of heat given off when a cer- 
tain weight of the substance combines with oxvgen. 

It will be found that not only is the heat of combustion 
a fixed quantity whether the union with oxygen takes place 
slowly or rapidly, but that the heat evolved in any given 
chemical reaction is always the same, and that chemical 
combination is always accompanied by an evolution of heat. 

Heat of Decomposition. — Just as it is true that a definite 
quantity of heat is evolved when two or more elements 
combine chemically, so also it is true that in order to de- 
compose the compound formed the same quantity of heat is 
absorbed. 



36 introduction to cBMismr. 

Chemical Energy and Chemical Work. — Any substance 
which has the power to unite with others can do chemical 
tvorlc, — it possesses chemical energy. Thus, all combustible 
substances can do work. In uniting with oxygen heat is 
evolved, and this can be transformed into motion. To go 
back to the example of the steam-engine, which was re- 
ferred to in an early part of the book, the cause of the 
motion is the burning of the fuel. We thus see that the 
source of the power in the steam-engine is chemical energy. 
Substances which have not the power to combine with 
others have no power to do chemical work, or they have no 
chemical energy. As far as power to combine with oxygen 
is concerned, water is a substance of this kind, as is also 
carbon dioxide, the gas formed when carbon is burned in 
oxygen. In order that they may do work, they must first 
be decomposed and their constituents put together in some 
form in which they have the power of combination. This 
decomposition of carbon dioxide and water is taking place 
constantly on the earth. All plant-life is dependent on 
it. The products of the action, i.e., the different kinds of 
wood, have energy, — they can do chemical work. This 
power to do work has been acquired from the heat of the 
sun, to which the decomposition of the carbon dioxide 
and water is mainly due. We have thus a transformation 
of the sun's heat into chemical energy, which is stored up 
in the combustible wood. The quantity of heat which 
would be given off in burning the wood would be exactly 
equal to the quantity used up in its formation. 

Oxides. — The compounds of oxygen with other elements 
are called oxides. To distinguish between different oxides 
the name of the element with which the oxygen is in com- 
bination is prefixed. Thus the compound of zinc and 
oxygen is called zinc oxide ; that of calcium and oxygen, 
calcium oxide; that of silver and oxygen, silver oxide, etc. 



CHAPTER III. 

HYDROGEN. 

In Experiment 4 it was found that when an electric 
current is passed through water two gases are obtained, 
one of which we have since studied and found to be oxy- 
gen. The other, it will be remembered, takes fire and 
burns, and is thus easily distinguished from oxygen. This 
second gas is hydrogen. 

Occurrence. — Hydrogen is found in nature very widely 
distributed, and in large quantity. It forms one ninth the 
weight of water, and is contained in all substances which 
enter into the composition of plants and animals. 

Preparation of Hydrogen. — It can be prepared : 

(a) By decomposition of water by means of the electric 
current ; 

(b) By decomposition of water by the action of certain 
metals ; 

(c) By the action of substances known as acids on metals. 
The following experiments will illustrate these methods: 

Experiment 26. — Eepeat Experiment 4 and examine the gases. 
Experiment 27. — Throw a small piece of sodium* on water. 

* The metals sodium and potassium are kept under oil. When a 
small piece is wanted take out one of the larger pieces from the 
bottle, roughly wipe off the oil with filter-paper aud cut off a 
piece the size needed. It is not advisable to use a piece larger than 
a pea. 

37 



38 INTRODUCTION TO CHEMISTRY. 

While it is floating on the surface apply a lighted match to 
it. A yellow flame will appear. This is burning hydrogen, the 
flame being colored yellow by the presence of the sodium, some of 
which also burns. Make the same experiment with potassium. 
The flame appears in this case without the aid of the match. It 
has a violet color which is due to the burning of some of the po- 
tassium. The gas given off in these experiments is either burned 
at once or escapes into the air. In the case of the potassium the 
action takes place rapidly, and the heat evolved is sufficient to set 
tire to the gas. In the case of the sodium the heat evolved does 
not set fire to the gas. In order to collect it unburned, it is only 
necessary to allow the decomposition to take place, so that the 
gas will rise in an inverted vessel filled with water. For this 
purpose fill a good-sized test-tube with water and invert it in a 
vessel of water. Cut off a piece of sodium not larger than a pea, 
wrap it in a layer or two of filter-paper, and with the fingers or 
a pair of curved forceps bring it quickly below the mouth of the 
test-tube and let it go. It will rise to the top, the decompo- 
sition of the water will take place quietly, and the gas formed, 
being unable to escape, will remain in the tube. By repeating 
this operation in the same tube a second portion of gas may be 
made, and so on until enough has been collected. 

Examine the gas and see whether it acts like the hydrogen 
obtained from water by means of the electric current. What 
evidence have you that they are the same ? Is this evidence suf- 
ficient to prove the identity of the two ? Examine the water on 
which the sodium or potassium has acted. Wet the fingers with 
it and rub them together. Taste the water. Does it change the 
color of red litmus paper ? 

Action of Sodium and Potassium on Water. — The explana- 
tion of the action of sodium and potassium on water will be 
given later. Suffice it for the present to say that water 
consists of hydrogen and oxygen, and that when sodium 
comes in contact with it this element takes the place of 
some of the hydrogen, forming the compound sodium hy- 
droxide or caustic soda. The action of potassium is of the 
same kind. The product is potassium hydroxide or caustic 
potash. 



DECOMPOSITION OF WATER. 



39 



Decomposition of Water by Iron.— Some metals which do 
not decompose water at ordinary temperatures, or which 
decompose it slowly, do so easily at elevated temperatures. 
This is true of iron. If steam is passed through a 
tube containing pieces of iron heated to redness, decompo- 
sition of the water takes place, the oxygen is retained 
by the iron, which enters into combination with it, while 
the hydrogen is liberated. 

Experiment 28. — In this experiment a porcelain tube with an 
internal diameter of from 20 to 25 mm. (about an inch) and a 
gas-furnace are desirable, though a hard-glass tube and a char- 
coal-furnace will answer. The arrangement of the apparatus is 
shown in Fig. 11. The hydrogen can be collected by displace- 




FlG. 11. 



ment of water, as in the case of oxygen. The products formed 
iire magnetic oxide of iron and hydrogen. 

Decomposition of Water by Carbon or Charcoal. — Many 

other substances have the power to decompose water and 
set hydrogen free. The fact that a combustible gas can 
be obtained from water has led to many attempts to manu- 
facture gas for heating and illuminating purposes from 
this substance. There is, however, no cheap substance 
known to us which has the power to decompose water at 
ordinary temperatures, All practicable methods involve 



40 INTRODUCTION TO CHEMISTRY. 

the use of heat, and it is not unfrequently the case that 
the quantity of heat required to effect the decomposition 
is greater than that which would be obtained by burning 
the hydrogen formed. In the manufacture of the so-called 
" water gas " which is now extensively used in the United 
States both for illuminating and heating purposes, water 
is decomposed by means of carbon which is used in the 
form of hard coal. Two gaseous products are formed both 
of which burn. They are carbon monoxide, or carbonic 
oxide, and hydrogen. This subject will be more fully dis- 
cussed under the head of carbon monoxide. 

Action of Acids upon Metals. — By far the most con- 
venient method for making hydrogen consists in treating 
a metal with an acid. As will be seen later, acids are 
substances which contain hydrogen, and which are char- 
acterized by the fact that they give up this hydrogen very 
easily and take up other elements in the place of it. 
Among the common acids fouud in every laboratory are 
hydrochloric acid, sulphuric acid, and nitric acid. The 
chemistry of these compounds will be taken up in due 
time; but as we shall be obliged to use them before they 
are studied systematically, a few words in regard to them 
are desirable at this time. 

Hydrochloric acid is a compound of hydrogen and chlo- 
rine. It is a gas which dissolves easily in water. It is 
this solution which we use in the laboratory, and which is 
manufactured in enormous quantities in connection with 
the manufacture of soda or sodium carbonate. It is fre- 
quently called " muriatic acid." 

Sulphuric acid is a compound of sulphur, oxygen, and 
hydrogen. It is an oily liquid and is frequently called 
" oil of vitriol." It is manufactured in very large quanti- 
ties, as it plays an important part in many of the most 
important chemical industries. 

Nitric acid is a compound containing nitrogen, oxygen, 






ACTION OF ACIDS UPON METALS. 



41 



and hydrogen. AY hen pure, it is a colorless liquid, though 
as we get it it is commonly colored somewhat yellow. 

When a metal, as zinc, is brought in contact with 
hydrochloric or sulphuric acid, an evolution of gas takes 
place at once. 

Experiment 20.— In a cylinder or test-tube put some small 
pieces of zinc, and pour upon it some ordinary hydrochloric acid. 
After the action has continued for a minute or two apply a 
lighted match to the mouth of the vessel. The gas will take fire 
and burn. If sulphuric acid diluted with five or six times its 
volume of water * is used instead of hydrochloric acid, the same 





Fig. 12. 



Fig. 13. 



result will be reached. The gas evolved is hydrogen. For the 
purpose of collecting the gas the operation is best performed in a 
bottle with two necks called a Wolff's flask (see Fig. 12), or in a 



* If it is desired to dilute ordinary concentrated sulphuric acid 
with water, the acid should be poured slowly into the water while 
the mixture is constantly stirred. If the water is poured into the 
acid, the heat evolved at the place where the two come in contact 
may be so great as to convert the water into steam and cause the 
strong acid to spatter. 



42 INTRODUCTION TO CHEMISTRY. 

wide-mouthed bottle in which is fitted a cork with two holes (see 
Fig. 13). Through one of the holes passes a funnel-tube, and 
through the other a glass tube bent in a convenient form. 

The zinc used is usually what is known as granulated zinc. 
It is prepared by melting it in a ladle and pouring the molten 
metal from an elevation of four or five feet into water. The 
advantage of this form is that it presents a large surface to the 
action of the acids. A handful of this zinc is introduced into 
the bottle, and enough of a cooled mixture of sulphuric acid and 
water (1 volume concentrated acid to 6 volumes water) poured 
upon it to cover it. Usually a brisk evolution of gas takes place 
at once. Wait for two or three minutes, and then collect some 
of the gas by displacement of water. When the action becomes 
slow, add more of the dilute acid. It will be well to fill several 
cylinders and bottles with the gas, and also a gasometer, from 
which it can be taken as it is needed for experiments. 

When zinc acts upon hydrochloric acid it takes the place of the 
hydrogen in the hydrochloric acid and forms the compound zinc 
chloride : 

Zinc + Hydrochloric Acid = Zinc Chloride -f Hydrogen ; 

or „. , Hydrogen ) Zinc J TT , 

ZmC + Chlorine S = Chlorine \ + *?*">&*■ 

When zinc acts upon sulphuric acid, it takes the place of the 
hydrogen and forms the compound zinc sulphate : 

Zinc + Sulphuric Acid = Zinc Sulphate + Hydrogen ; 

or Sulphur ] Sulphur \ 

Zinc + Oxygen V = Oxygen > + Hydrogen. 
Hydrogen ) Zinc ) 

Experiment 30. — After the action is over pour the contents of 
the flask through a filter into an evaporating-dish, and boil off 
the greater part of the water, so that, on cooling, the substance 
contained in solution will be deposited. If the operation is car- 
ried on properly, the substance will be deposited in regular forms 
called crystals. It is zinc sulphate. 

Physical Properties of Hydrogen. — Hydrogen is a color- 
less, inodorous, tasteless gas. Made by the action of zinc 



PHYSICAL PROPERTIES OF HYDROGEN. 



43 



on acids, it has a slightly disagreeable odor. This is due 
to the presence of impurities. If it is passed through 
certain substances which have the power to destroy the 
impurities, the odor is destroyed. 

Experiment 31. — Pass some of the gas through a wash-cylinder 
containing a solution of potassium permanganate; collect some 
of it, and notice whether it has an odor or not. The apparatus 
should be arranged as shown in Fig. 14. The solution of potas- 




Fig. 14. 



sium permanganate is, of course, contained in the small cylinder 
.4, and the tubes are so arranged that the gas bubbles through it. 

The gas is not poisonous, and may therefore be inhaled 
with impunity. We could not, however, live in an atmos- 
phere of hydrogen, as we need oxygen. It is the lightest 
known substance, being fourteen and a half times lighter 
than the air and sixteen times lighter than oxygen. Its 
lightness may be shown by a number of simple experi- 
ments. 



44 INTRODUCTION TO CBEMISTHY. 

Experiment 32. — Place a vessel containing hydrogen with the 
mouth upward and uncovered. In a short time examine the gas, 
and see whether it is hydrogen. 

Experiment 33. — Gradually bring a vessel containing hydrogen 
with its mouth upward below an inverted vessel containing air, 
in the way shown in Fig. 15. The air will be displaced. On ex- 




Fig. 15. 



animation the inverted vessel will be found to contain hydrogen, 
while the one with the mouth upward will contain none. The 
gas is thus poured upwards. 

Experiment 34. — Soap-bubbles filled with hydrogen rise in 
the air. This experiment is best performed by connecting an 
ordinary clay pipe by means of a piece of rubber tubing with the 
exit-tube of a gasometer filled with hydrogen. Small balloons of 
collodion are also made for the purpose of showing the lightness 
of hydrogen. 

Balloons are always filled with hydrogen, or some other 
light gas. Some kinds of illuminating gas are rich in hy- 
drogen, and may therefore be used for the purpose. 

A litre of hydrogen at 0° (Centigrade), and under the 
pressure of 760 mm., weighs 0.089578 gram. Its specific 
gravity is 0.0691. A comparison of these figures with the 
corresponding figures for oxygen leads to an interesting 
observation. The weight of a litre of oxygen is 1.429 
grams; its specific gravity is 1.10563. The ratio of the 



CHEMICAL PROPERTIES OF HYDROGEN. 



45 



weight of equal volumes of hydrogen and oxygen to each 
other is 1 : 16, or 

0.089578 : 1.429 : : 1 : 16. 



Although it has been subjected to a very low tempera- 
ture and high pressure, it is doubtful whether hydrogen 
has ever been liquefied. 

Chemical Properties of Hydrogen. — Under ordinary cir- 
cumstances, hydrogen is not a particularly active element. 
It does not unite with oxygen at ordinary temperatures, 
but, like w r ood and most other combustible substances, 
needs to be heated to the kindling temperature before it 
will burn. We have seen that it burns when a lighted match 
is applied to it. The flame is colorless, or very slightly 
blue. As burned under ordinary circumstances, the flame 
is colored, in consequence of the presence of foreign sub- 
stances; but that it is colorless when the gas is burned 
alone can be shown by burning it from a platinum tube, 
which is itself not acted upon by the heat. 

Experiment 35. — If there is no small platinum tube available, 
roll up a small piece of platinum foil and melt it into the end of 
a glass tube, as shown in Fig. 16. Connect the 
burner thus made with the gasometer containing 
hydrogen, and after the gas has been allowed to 
issue from it for a moment, set fire to it. In a 
short time it will be seen that the flame is practi- 
cally colorless, and gives no light. That it is hot 
can readily be shown by holding a piece of plati- 
num wire or a piece of some other metal in it. 

Hydrogen burns. We have seen that this 
consists in combining with oxygen. On the 
other hand, substances which burn in the air FlG - 16 - 
are extinguished when put into a vessel containing hydro- 
gen. This is equivalent to saying that a substance which 




46 



INTRODUCTION TO CHEMISTRY. 



is uniting with oxygen does not continue to unite with 
oxygen when put in an atmosphere of hy- 
drogen, and does not combine with the 
hydrogen. This is expressed by saying 
that hydrogen does not support combus- 
tion. The following experiment shows 
this. 

Experiment 36.— Hold a cylinder filled with 
hydrogen with the mouth downward. Insert 
into the vessel a lighted taper held on a bent 
wire, as shown in Fig. 17. The gas takes fire 
at the mouth of the vessel, but the taper is 
extinguished. On withdrawing the taper and 
holding the wick for a moment in the burning 
hydrogen, it will take fire, but on putting it 
back in the hydrogen it will again be extin- 
guished. Other burning substances should 
be tried in a similar way. 




Fig. IT 



Product Formed when Hydrogen Burns in Oxygen. — As 

when hydrogen burns it combines with oxygen, a product 
should be obtained in which both hydrogen and oxygen 
are present. In the experiments performed we have seen 
no evidence of the formation of such a product, simply 
for the reason that when formed it is an invisible gas, and, 
though it can easily be condensed to a liquid, no precau- 
tions were taken to get it in this form. The product is, 
in fact, ordinary water, which we shall next study. 



CHAPTER IV. 

COMBINATION OF HYDROGEN AND OXYGEN.— WATER. 

Water was regarded as an element until, towards the 
end of the last century, the discovery of hydrogen and 
oxygen, and of the nature of combustion, led to the dis- 
covery of its true composition. 

Occurrence. — The wide distribution of water on the earth 
is familiar to every one. But water also occurs in forms 
and conditions which prevent its immediate recognition. 
Thus all living things contain a large proportion of water, 
which can be driven off by heat. If a piece of wood or a 
piece of meat is heated, water passes off. 

Experiment 37. — In a dry tube heat gently a small piece of 
wood. What evidence do you obtain that water is given off ? Do 
the same thing with a piece of fresh meat. 

The proportion of water in animal and vegetable sub- 
stances is very great. If the body of a man weighing 150 
pounds should be put in an oven and thoroughly dried, 
there would be left only about 40 pounds of solid matter, 
all the rest being water. As all meat, vegetables, and food- 
stuffs in general contain a similar large proportion of water, 
it is evident that water is a very important article of com- 
merce. When we buy four pounds of beef, we pay for 
about three pounds of water and one of solid matter. 

Water of Crystallization. — Water also occurs in another 
form in which it does not easily reveal its presence. This 
is as ivater of crystallization. 

47 



48 INTRODUCTION TO CHEMISTRY. 

Experiment 38. — Take some of the crystals of zinc sulphate 
obtained in Experiment 30. Spread them out on a layer of filter- 
paper, and finally press two or three of them between folds of 
the paper. Examine them carefully. They appear to be quite 
dry, and in the ordinary sense they are dry. Put them into a 
dry tube, and heat them gently. What evidence do you obtain of 
the giving off of water ? Describe the changes which the crystals 
undergo. 

Experiment 39. — Perform a similar experiment with some 
gypsum, which is the natural substance from w r hich " plaster of 
Paris " is made. 

Experiment 40. — Heat a few small crystals of copper sulphate 
or blue vitriol. What evidence of water ? Describe the changes 
in the crystals. After no further change takes place, dissolve 
what is left in a little water. What is the color of the solution ? 
Evaporate to the point of crystallization. How do the crystals 
obtained compare with those first taken ? 

Many compounds when deposited from solutions in water 
in the form of crystals combine with definite quantities of 
water. This w r ater is not present as such, but is held in 
chemical combination. Hence the substance does not ap- 
pear moist, though it may contain more than half its weight 
of water. This water of crystallization is, in some way 
which we do not understand, essential to the form of the 
crystal. If it is driven off by heat, the crystal is destroyed. 
Some compounds combine under different circumstances 
with different quantities of water, the form of the crystals 
varying with the quantity of water in combination. 

Efflorescence. — Compounds differ greatly as regards the 
ease with which they give up water of crystallization. In 
general, it is given off when the compound containing it is 
heated to the temperature of boiling water. But some 
compounds give it up by simple contact with the air. 
This is true of sodium sulphate, or Glauber's salt, which 
contains about 56 per cent of water of crystallization. 

Experiment 41. — Select a few crystals of sodium sulphate 
which have no' lost their lustre. Put them on a watch-glass, 



DELIQUESCENCE— ANALYSIS AND SYNTHESIS. 49 

and let them lie exposed to the air for an hour or two. What 
change takes place in their appearance ? How does this change 
compare with that of the crystals of zinc sulphate ? 

Compounds which lose their water of crystallization by 
simple contact with the air are said to effloresce. They are 
called efflorescent. 

Deliquescence. — Some compounds if deprived of their 
water of crystallization will take it up again when allowed 
to lie in an atmosphere containing moisture. As the air 
always contains moisture, it is only necessary to expose 
such compounds to the air in order to notice the phenome- 
non. It is well shown by the compound calcium chloride. 
This substance has a remarkable power of attracting water 
to itself and holding it in combination. 

Experiment 42. — Expose a few pieces of calcium chloride to 
the air for some hours. Describe the changes that take place. 

Substances which absorb water from the air are said to 
deliquesce. They are called deliquescent. 

Analysis and Synthesis. — In order to determine the com- 
position of water, as well as that of any other compound, 
we must analyze it. We may simply determine what sub- 
stances enter into its composition without determining the 
relative quantities of these substances. In this case we 
make what is called a qualitative analysis. If, however, 
we not only determine what substances are present, but 
also in what quantities they are present, we then make a 
quantitative analysis. 

The composition of a substance may also be determined 
by putting together its constituents and causing them to 
combine chemically. An operation of this kind is called a 
synthesis. A synthesis, then, is the opposite of an analy- 
sis. Just as w r e may make a qualitative or a quantitative 



50 INTRODUCTION TO CHEMISTRY. 

analysis, so also we may make a qualitative or a quantita- 
tive synthesis. These processes are well illustrated in the 
operations necessary to determine the composition of water. 

Decomposition of Water by the Electric Current and 
what it Teaches. — That water contains hydrogen and 
oxygen has already been shown in Experiment 4. It will 
now be w T ell to repeat the experiment and see whether we 
can learn anything more regarding the composition of w T ater 
than that it contains hydrogen and oxygen. In the first 
place, the question suggests itself, In what proportions, by 
weight and by volume, are the gases combined ? 

Experiment 43. — The tubes in the apparatus used in Experi- 
ment 4, or some other similar apparatus, should be marked by 
means of a file, or by etching, so that equal divisions can be 
recognized. Tubes thus divided so that the divisions indicate 
cubic centimetres are most convenient for the purpose. Let the 
gases formed by the action of the electric current, as in Experi- 
ment 4, rise in the graduated tubes, and observe the volumes. 
It will be seen that when one tube is just full of gas, the other, if 
it is of the same size, will be only half full. On examining the 
gases the larger volume will be found to be hydrogen, and the 
smaller volume oxygen. This experiment has been performed an 
untold number of times and always with the same result. 

The relative weights of equal volumes of the two gases are 
known, so that the relative weights of the gases obtained 
from water by the action of the electric current can easily be 
calculated. The ratio of the weights of equal volumes of 
hydrogen and oxygen is 1 : 16. Therefore, if two volumes of 
hydrogen are combined with one volume of oxygen, the 
ratio between the weights is 2 : 16 or 1 : 8. Although the 
above experiment shows that hydrogen and oxygen are 
obtained from water in the proportion of two volumes of 
the former to one of the latter, or of one part by weight 
of the former to eight parts by weight of the latter, th§ 



SYNTHESIS OF WATER BY BURNING HYDROGEN 51 

experiment does not prove that this is the actual compo- 
sition of water. For it may be that other elements besides 
hydrogen and oxygen are contained in the water, and it 
may be that all the hydrogen and oxygen are not set free 
by the action of the electric current. Whether either of 
these possibilities is true might be determined by decom- 
posing a weighed quantity of water, and weighing the 
hydrogen and oxygen obtained from it. If it should be 
found that the sum of the weights of hydrogen and oxy- 
gen is equal to the weight of the water decomposed, this 
fact would be evidence that only hydrogen and oxygen are 
contained in water, and that they are present in the pro- 
portions stated. 

Synthesis of Water by Burning Hydrogen. — That water 
consists of hydrogen and oxygen only can be satisfactorily 
proved by effecting its synthesis. In the first place, it 
can be shown that water is formed when hydrogen burns 
in the air, and, as it has been shown that burning is com- 
bining with oxygen, the conclusion is justified that water 
consists of hydrogen and oxygen. 

Experiment 44. — Pass hydrogen from a generating-flask or a 
gasometer through a tube containing some substance that will 
absorb moisture, for all gases made in the ordinary way and col- 
lected over water are charged with moisture. We have seen in 
Experiment 42 that calcium chloride has the power to absorb 
moisture. It is extensively used in the laboratory for the purpose 
of drying gases, and it may be used in the present experiment. 
It should be in granulated form, not powdered. After passing 
the hydrogen through the calcium chloride, pass it through a 
tube ending in a narrow opening, and set fire to it. If now a dry 
vessel be held over the flame, drops of water will condense on its 
surface and run down. A convenient arrangement of the appa- 
ratus is shown in Fig. 18. 

A is the calcium-chloride tube. Before lighting the jet hold a 
glass plate in the escaping gas, and see whether water is deposited 
on it. Light the jet before putting it under the bell jar, other- 



52 



INTRODUCTION TO CHEMISTRY. 



wise if hydrogen is allowed to escape into the vessel it will con- 
tain a mixture of air and hydrogen, and this mixture, as will soon 
be seen, is explosive. 




Fig. 18. 

Synthesis of Water by Mixing Hydrogen and Oxygen. — 

If hydrogen and oxygen are mixed together, and the mix- 
ture is allowed to stand unmolested, it remains unchanged c 
If, however, a spark or a flame is brought in contact with 
the mixture, a violent explosion occurs, and a careful ex- 
amination has shown that the explosion is the result of 
the combination of the two gases. The noise is caused by 
the sudden great expansion of the gases occasioned by the 
development of heat. This expansion is instantly followed 
by a contraction. 

Experiment 45. — Mix hydrogen and oxygen in the proportion 
of about 2 volumes of hydrogen to 1 volume of oxygen in a gas- 
ometer. Fill soap-bubbles, as directed in Experiment 34, with 
this mixture and allow them to rise in the air. As they rise 
bring alighted taper in contact with them, when a sharp explosion 
will occur. Great care must be taken to keep all flames* away 
from the vicinity of the gasometer and the end of the delivery-tube. 



QUANTITA11VE SYNTHESIS OF WATER 53 

This experiment may be conveniently performed by hanging 
up, about six to eight feet above the experiment-table, a good- 
sized tin funnel-shaped vessel with the mouth downward. Now 
place a gas jet or a small flame of any kind at the mouth of the 
vessel. If the soap-bubbles are allowed to rise below this appara- 
tus they will come in contact with flame and explode at once * 

This experiment simply shows that a mixture of hydro- 
gen and oxygen explodes when brought in contact with a 
flame, and that the gases do not act upon each other at 
ordinary temperatures. 

Quantitative Synthesis of Water. — In order to show that 
when the explosion occurs water is formed, and in what 
proportions the gases combine, it is necessary to w r ork in 
closed vessels so constructed as to permit accurate measure- 
ment of the volumes of the gases. The experiment is so 
important that, if possible, it should be performed, at least 
by the teacher, before the class. The vessel in which the 
gases are brought together and caused to combine is called 
a eudiometer (from evSia, calm air, and juerpov, a 
measure). It is simply a tube graduated in millimetres 
and having two small platinum wires passed through it at 
the closed end, nearly meeting inside and ending in loops 
outside, as shown in Fig. 19. The eudiometer is filled with 



Fig. 19. 

mercury, inverted in a mercury trough, and held in an up- 
right position by means of proper clamps. A quantity of 
pure hydrogen is passed up into the tube and its volume 
accurately measured. About half thi& volume of oxygen 
is then introduced and accurately measured, and after the 

* The same apparatus may be used in experimenting with soap- 
bubbles filled with hydrogen. 



54 INTRODUCTION TO CHEMISTRY. 

mixture has been allowed to stand for a few minutes, a 
spark is passed between the wires in the eudiometer by- 
connecting the loops with the poles of a small Ruhmkorff 
coil or with a Leyden jar. Under these circumstances the 
explosion takes place noiselessly and with very little danger. 
If the interior of the tube was dry before the explosion, it 
will be seen to be moist afterwards, and a marked decrease 
in the volume of the gases is also observed. That water is 
the product of the action has been proved beyond any 
possibility of a doubt, over and over again. As the liquid 
water which is formed occupies an almost inappreciable 
volume as compared with the volume of the gases which 
combine, the decrease in volume represents the total 
volume of hydrogen and oxygen which have combined. 
Now, if the experiment is performed with the two gases 
in different proportions, it will be found that only when 
they are mixed in the proportion of two volumes of hydro- 
gen and one volume of oxygen do they completely dis- 
appear in the explosion. If hydrogen is present in larger 
proportion, the excess is left over. If oxygen is present 
in larger proportion, the excess of oxygen is left over. It 
appears, therefore, that when hydrogen and oxygen com- 
bine to form water, they do so in the proportion of two 
volumes of hydrogen to one volume of oxygen. In order 
that the student may fully appreciate this experiment, it 
is desirable that he should at this point familiarize himself 
with the precautions necessary in measuring the volumes 
of gases, if he has not already done so. 

Correction for Temperature. — The volume of a gas varies 
with the temperature and pressure. When the temperature 
of a gas is raised one degree Centigrade its volume is in- 
creased j\j part. If, therefore, the volume of a gas at 0° 
is V 9 at t° its volume v will be 



'9 



F+^.F or v = V+±- r V. 



TEMPERATURE- AND PRESSURE-CORRECTIONS. 55 
This expression may also be written 

v = V+ 0.00366/ . V or v = V(l -f 0.00366/). 
From this it follows that 

V= v - . 

1 + 0.00366/ 

It is customary to reduce the observed volume of a gas 
to the volume which it would have at 0°. The correction 
is made in accordance with the above expression. Thus, 
if the volume of a gas is found to be 250 cubic centimetres 
at 15°, and it is required to find what its volume would be 
at 0°, the calculation is made thus: In this case v, the ob- 
served volume, is 250 cc; /, the temperature, is 15°. Sub- 

v 
stituting these values in the equation V = , 

1 ~\~ U.UUooOf 

we have 

V= ^__ 

1 + 0.00366.15' 

from which we get 236.99 as the value of F. 

But the volume of a gas varies also according to the 
pressure. When the pressure is doubled, the volume is de- 
creased to one half; and when the pressure is decreased to 
one half, the volume is doubled : and so on. In other words, 
the volume of a gas varies inversely with the pressure. 
Increase the pressure two, three, or four times and the 
volume becomes one half, one third, or one fourth, and 
vice versa. If the gas has the volume Fat the pressure P, 
and at pressure p the volume r, these values bear to one 
another the relations expressed in the equation 

PV=pv. 

Correction for Pressure. — The pressure is usually stated 
in millimetres, and reference is to the height of a column 
of mercury to which the pressure corresponds. A gas 
contained in an open vessel, or in a vessel over mercury or 



56 



INTRODUCTION TO CHEMISTRY. 



water, in which the level of the liquid inside and outside 
the vessel is the same, is under the pressure of the atmos- 
phere. What that is we learn from the barometer. As 
this pressure varies, it is necessary to read the barometer 
whenever a gas is measured, and then to reduce the ob- 
served volume to certain conditions which are accepted as 
standard. When the gas is measured in a tube over mercury 
or water, and the level of the liquid inside the tube is 
higher than that outside, the gas is under diminished 
pressure, the amount of diminution depending on the 
height of the column of mercury or water in the tube. 

Thus, if the arrangement is as 
represented in Fig. 20, and the 
height of the mercury column 
above the level of the mercury in 
the trough is 100 millimetres, and 
the pressure of the atmosphere is 
760 millimetres, then the gas in 
the tube is not under the full at- 
mospheric pressure, for the atmos- 
pheric pressure exerted on the 
mercury is supporting a column of 
mercury 100 millimetres high, and 
the pressure actually brought to 
bear on the gas corresponds to 
760 - 100 = 660 mm. Suppose 
that in this case the volume of gas actually measured is 
75 cc. Call this v. What would be the actual volume 
V of the gas under the standard pressure 760 mm. i We 
have seen that 

VP = vp. 
Now P - 760, v = 75, and p = 660. Therefore, 

760 V = 75 X 660, or 




Fig. 20. 



7 5 X 660 
V — — ^^ — = 65.13. 



760. 



CORRECTION FOR WATER-VAPOR. 57 

Combined Volumetric Corrections. — In all cases it is neces- 
sary to make a correction similar to this in dealing with 
the volumes of gases. The correction for temperature 
and that for pressure may be made in one operation, the 
formula being 

vp 



V = 



760(1 + 0.00366*)' 



in which F=the volume of the gas at 0° and 760 mm. 
pressure; v = the observed volume; t — the observed tem- 
perature; p = the pressure under which the gas is meas- 
ured. 

Correction for Aqueous Tension. — The presence of water- 
vapor in a gas also influences its volume, and this must be 
taken into account. The formula for making all the cor- 
rections required in determining the volume of a gas is 

r= V (P ~ a ) 



760(1 + 0.003660' 
in which a is the tension of water-vapor at t°. 

[Problems. — The volume of a gas contained in a eudiometer 
measures 42 cc. The height of the mercury column over which 
it stands is 68 mm. The barometer indicates an atmospheric 
pressure of 746 mm. The temperature is 18° C. What would be 
the volume of the gas at 0° and 760 mm. pressure ? 

The volume of a gas contained in a vessel over a column of 
mercury 85 mm. high measures 24 cc. The barometer indicates 
a pressure of 774 mm. The temperature is 19°. What would be 
the volume of the gas under normal conditions, i.e. K t = 0° and 
P =• 760 mm. ? 

The volume of a gas contained in a vessel over a liquid, the 
level of the liquid inside and outside being the same : v = 80 cc. ; 
t = 20° ; p = 740 mm. What is the value of For the volume at 
0° and 760 mm. ?] 

Apparatus for Measuring the Volume of a Gas. — A conven- 
ient apparatus for measuring gas-volumes is that represented 



58 



INTRODUCTION- TO CHEMISTRY. 




in Fig. 21. It consists of two tubes 
connected at the base by means of a 
piece of rubber tubing and contain- 
ing water. The tube A is graduated, 
the other is not. The gas to be 
measured is brought into the tube A, 
and the other tube is then placed at 
the side of the one containing the 
gas, and its height adjusted so that 
the column of liquid in both tubes 
is at the same level. Under these 
circumstances, obviously the gas is 
under the atmospheric pressure for 
which the necessary correction must, 
of course, be made. It is also neces- 
sary in this case to make the cor- 
rections for temperature and for 
the tension of water-vapor. It is, 




Fig. 21. 



Fig. 22. 



CALCULATION OF ftfiSULTS. 59 

further, sometimes convenient when the gas is measured 
over water to transfer the measuring-tube to a vessel con- 
taining enough water to permit the immersion of the tube 
to a point at which the level of the liquid inside and out- 
side the tube is the same. In this case the conditions 
are the same as in the apparatus just described. The ar- 
rangement is shown in Fig. 22. 

Calculation of the Results Obtained on Exploding Mix- 
tures of Hydrogen and Oxygen. — Having determined that 
whenever hydrogen and oxygen combine, they do so in the 
proportion 1 volume of oxygen to 2 volumes of hydrogen, 
and that when they combine, the volume of water formed 
is so slight as to amount to nothing in the measure- 
ments, we know that whenever a mixture of hydrogen 
and oxygen is exploded, no matter in what proportions 
they may be present, the volume of gas which disappears 
as such consisted of 2 volumes of hydrogen and 1 volume 
of oxygen, or, in other words, one third of the volume 
which disappears was oxygen and two thirds hydrogen. 
Take this example : A quantity of hydrogen corresponding 
to 60 cc. under standard conditions is introduced into a 
eudiometer; 40 cc. of oxygen are added. What contrac- 
tion will there be on exploding the mixture ? Plainly the 
60 cc. of hydrogen will combine with 30 cc. of oxygen. 
The 90 cc. of gas will disappear, and the 10 cc. of oxygen 
will remain unchanged. From a total volume of 100 cc, 
therefore, we get a contraction to 10 cc. One third of the 
contraction represents the oxygen and two thirds the 
hydrogen. 

Synthesis of Water by Passing Hydrogen over Heated 
Oxides. — The synthesis of water can be effected by passing 
hydrogen over a compound containing oxygen and heated 
to a sufficiently high temperature. A convenient sub- 
stance for this purpose is the compound of copper and 



60 



INTRODUCTION TO CHEMISTRY. 



oxygen known as copper oxide or black oxide of copper. 
When hydrogen is passed over this compound at ordinary 
temperatures no action takes place. If, however, the tem- 
perature is raised to low redness the hydrogen combines 
with the oxygen, forming water, and the copper is left 
behind as such. 

Experiment 46.— Arrange an apparatus as shown in Fig. 23. 
A is a Wolff's flask for generating hydrogen. To remove im- 
purities the gas is passed through a solution of potassium per- 




Fig. 23. 

manganate contained in the wash-cylinder B. The cylinder C 
contains concentrated sulphuric acid, and the U-shaped tube 
D contains granulated calcium chloride, both of them serving to 
remove moisture from the gas. The pure dry hydrogen is now 
passed through the hard-glass tube 2£, which contains a layer of 
copper oxide. After the apparatus is filled with hydrogen the 
gas jet is lighted and the copper oxide heated to low redness. 
What evidence do you obtain of the formation of water ? What 
change takes place in the color of the substance in the tube ? Try 
the action of nitric acid on a little of the black oxide of copper, 
and on the substance left after the action of the hydrogen. 

Quantitative Synthesis of Water. — In this case the loss 
in weight of the copper oxide represents oxygen. If, 
therefore, we should weigh the copper oxide before the 



QUANTITATIVE SYNTHESIS OF WATER. 61 

experiment, and afterwards the copper, and should also 
collect and weigh the water formed, we could from the 
figures obtained easily calculate the relative weight of the 
oxygen contained in the water. The water can easily be 
collected by passing it into a tube filled with calcium chlo- 
ride. If the tube is weighed before the experiment and 
after it, the gain in weight will represent the weight of the 
water collected. All these weighings can be made without 
difficulty on a chemical balance such as is found in every 
chemical laboratory. Where time will permit it will be 
well for each student to go through with this experiment. 
A few experiments of this kind w T ill serve to impress upon 
the mind the reality of the quantitative relations about 
which he is constantly hearing. If it is performed, a small 
hard-glass tube from 12 to 15 centimetres (5 to 6 inches) 
long and about 1 centimetre (or half an inch) internal 
diameter should be used in place of the tube E in the 
qualitative experiment above described. The tube is 
drawn out at one end and a small plug of asbestos put in 
the small end. Connection with the weighed calcium- 
chloride tube is made at this end. The tube is first dried 
thoroughly and weighed. Then a few grams of coarsely 
granulated copper oxide are introduced into it. After the 
experiment the tube and the copper are weighed again. 
The calcium-chloride tube should of course be weighed 
before and after the experiment. The results are cal- 
culated thus : 

Let x = weight of tube + copper oxide before the 

experiment ; 
y = weight of tube + copper after the experi- 
ment. 

Then x — y = weight of oxygen taken from the copper 
oxide. 

Let a = weight of calcium-chloride tube before; 

o = " " " " after, 

Then b — a — weight of water formed. 



62 INTRODUCTION TO CHEMISTRY. 

If the experiment is carefully performed, it will be found 
that the ratio *- — — is very nearly f . 

Oxidation and Reduction.— Any substance which like 
hydrogen has the power to abstract oxygen from com- 
pounds containing it, is called a reducing agent. The 
process of abstracting oxygen from a compound is called 
r eduction. Reduction and oxidation are therefore comple- 
mentary processes. We shall hereafter become acquainted 
with a number of important and interesting reducing 
processes. 

The Oxyhydrogen Blowpipe. — The heat evolved when 
hydrogen combines with oxygen is very great, and it is 
utilized for various purposes. To burn hydrogen in the 
air is, as we have seen, a simple matter, but to burn it in 
oxygen requires a special apparatus to prevent the mixing 
of the gases before they reach the end of the tube where 
the combustion takes place. The oxyhydrogen blowpipe 
answers this purpose. It is simply a tube with a smaller 
tube passing through it, as shown in Fig. 24. 







Fig. 24. 



The hydrogen is admitted through a and the oxygen 
through b. It will be seen that they come together only at 
the end of the tube. The hydrogen is first passed through 
and lighted ; then the oxygen is passed through slowly, the 
pressure being increased until the flame appears thin and 
straight. It gives very little light, but is intensely hot. 

Experiment 47. — Hold in the flame of the oxyhydrogen blow- 
pipe successively a piece of iron wire, a piece of a steel watch- 



THE LIME LIGHT— NATURAL WATERS. 6S 

spring, a piece of copper wire, a piece of zinc, a piece of platinum 
wire. 

The metal platinum is used for many purposes, particu 
larly for chemical operations. The vessels are made from 
molten platinum, and the metal is melted by means of the 
oxyhydrogen blowpipe. 

The Lime Light. — When this flame is allowed to play upon 
some substance which it cannot melt nor burn up, the sub- 
stance becomes heated so high that it gives off an intense 
light. The substance commonly used is quicklime. Hence 
the light is often called the lime light. It is also known as 
the Drummond light. 

Experiment 48. — Cut a piece of lime of convenient size and 
shape, say 25 mm. (1 inch) long by 20 mm. (f inch) wide, and 
the same thickness. Fix it in position so that the flame of the 
oxyhydrogen blowpipe will play upon it. The light is very bright, 
but by no means as intense as the electric light. 

Natural Waters. — The purest water found in nature is 
rain-water, particularly that which falls after it has rained 
for some time. That which first falls always contains im- 
purities from the air. As soon as the rain-water comes in 
contact with the earth, and begins its course tow 7 ards the 
ocean, it begins to take up various substances, according to 
the character of the soil with which it comes in contact. 
Mountain streams which flow over rocky beds, particularly 
over beds of sandstone, which is very insoluble, contain 
exceptionally pure water. Streams which flow over lime- 
stone dissolve some of the stone, and the water becomes 
" hard." The many varieties of mineral springs have their 
origin in the presence in the earth of certain substances 
which are soluble in w r ater. Common salt occurs in large 
quantities in different parts of the earth. As it is easily 
soluble in water, many streams contain it ; and as all the 
streams find their way into the ocean, we see one reason 
why the water of the ocean should be salt. As streams ap- 



64 



INTRODUCTION TO CHEMISTRY. 



proach the habitations of man they are subjected to a serious 
cause of contamination. The drainage from the neighbor- 
hood of human dwellings is very apt to find its way into a 
near stream. This condition of things is most strikingly 
illustrated by the case of a large town situated on the banks 
of a river. It frequently happens that the water of the 
river is used for drinking purposes, and it also frequently 
happens that the water is contaminated by drainage. Water 
when once contaminated by drainage tends to become pure 
again by contact with the air. If it is to be used for drink- 
ing purposes, however, it is not well to rely too much upon 
this process of purification. 

Distillation of Water. — In order to get pure water, it 
must be distilled. Distillation consists in boiling the water, 
and then condensing the vapor by passing it through a tube 
which is kept cool by surrounding it with cold water. A 
simple apparatus for the purpose is that illustrated in 
Fig. 25. 







Fig. 25. 



The water to be distilled is placed in the flask A. The 
flask is connected by means of a bent glass tube B with 
the long tube CO. This in turn is surrounded by the 
larger tube or jacket D. The side tube E is connected 



PROPERTIES OF WATER—WATER AS A SOLVENT. 65 

with a faucet by means of the rubber tube G. The water 
is allowed to flow slowly into the jacket and out at F, 
whence it passes through the rubber tube H to the sink. 
When the water in A is boiled, the vapor passes into the 
tube CO. Here it is cooled down, and takes the form of 
liquid, which runs down and collects in the flask K, which 
is called the receiver. 

Experiment 49. — Dissolve some copper sulphate, or other col- 
ored substance, in a litre of water, and distil the water. 

Properties of Water. — Pure water is tasteless and inodor- 
ous. In thin layers it is colorless, but in thick layers it is 
blue. This has been shown in the laboratory by filling a 
long tube with distilled water. When looked through it 
appears blue. The beautiful blue color of the water of 
some lakes is the natural color of pure water. 

On cooling, water contracts until it reaches the tempera- 
ture of 4°. At this point it has its maximum density. 
When cooled below 4° it expands, and the specific gravity 
of ice is somewhat less than that of water. Hence ice floats 
on water. 

Water as a Solvent. — It is known that many solids, 
liquids, and gases when brought into water disappear and 
form colorless liquids which look like water. Some give 
colored liquids of the same color as the substance dis- 
solved, and others give liquids which have colors quite dif- 
ferent from the dissolved substances. On the other hand 
there are many substances that do not dissolve in water. If 
a very small quantity of substance is dissolved in a large 
quantity of water, and the solution thoroughly stirred, the 
dissolved substance is uniformly distributed throughout the 
liquid, as can be shown by refined chemical methods. That 
the dissolved substance is everywhere present in the solu- 
tion can be shown, further, by the aid of certain dye- 
stuffs, as, for example, magenta. A drop of a concentrated 
solution of this substance brought into many gallons of 



66 INTRODUCTION TO CHEMISTRY. 

water imparts a distinct color to all parts of the liquid. An 
experiment of this kind gives some idea of the extent to 
which the subdivision of matter can be carried. For it is 
evident that in each drop of the dilute solution there must 
be contained some of the coloring matter, though the quan- 
tity must be what we should ordinarily speak of as infinites- 
imal. While there seems to be no limit to the extent to 
which a solution can be diluted, and still retain the dis- 
solved substance uniformly distributed through its mass, 
there is a limit to the amount of every substance that can 
be brought into solution, and this varies with the tempera- 
ture, and, in the case of gases, with the pressure. Some 
substances are easily soluble ; others are difficultly soluble. 
When the solutions are boiled the water simply passes off 
and leaves the dissolved substance behind, if it is a non- 
volatile solid. If, however, the substance in solution is a 
liquid a partial separation will take place, the extent of the 
separation depending largely upon the difference between 
the boiling-points of the water and the other liquid. If, 
finally, the substance in solution is a gas, it generally passes 
off when the solution is heated, though in some cases water 
is given off, leaving the gas in solution. 

Solutions and Chemical Compounds. — Solutions, in gen- 
eral, seem to differ from true chemical compounds in some 
important particulars, and also from mere mechanical mix- 
tures. Definiteness of composition is a common character- 
istic of chemical compounds, but solutions have no definite 
composition. Any quantity of a substance, from the mi- 
nutest particle to a certain fixed quantity, can be dissolved, 
and the solution formed is in each case uniform and ap- 
pears to be a perfect solution. The subject of solution is 
at present under investigation, and much light has been 
thrown upon it. 

Solution as an Aid to Chemical Action. — When it is de- 
sired to secure the chemical action of one solid substance 






OZONE-HYDROGEN DIOXIDE. 67 

upon another it is frequently found advantageous to bring 
them together in solution. The full explanation of this re- 
mains yet to be given, but it appears highly probable that 
when a substance is dissolved in water some deep-seated 
change takes place in it, and that this is one of the principal 
reasons why substances in solution act upon each other so 
readily. 

Ozone. — When electric sparks are passed for a time 
through oxygen the gas undergoes a remarkable change. 
It acquires a strong odor, and is much more active than 
under ordinary circumstances. The odor of the gas is 
observed in the neighborhood of an electric machine in 
action, and is said to be noticed during thunder-showers. 
The substance which has the odor is ozone. It is formed 
in a number of chemical reactions, as when phosphorus 
acts on the air in the presence of water. By cold and 
pressure it has been changed to a dark-blue liquid. 

AVhen a certain volume of oxygen is converted into 
ozone the volume of gas is decreased from three to two. 

By heating ozone above 300° it is converted into ordinary 
oxygen, and its volume is increased from two to three. 

It is clear that the element oxygen can be converted into 
something else without the addition of anything to it. 
This might lead us to conclude that it is not an element. 
But the substance formed from it has exactly the same 
weight and can be changed back again to oxygen without 
anything being added to it. It follows that the change 
must take place within the oxygen itself. The commonly 
accepted explanation of the relation between oxygen and 
ozone will be given later. 

Ozone is present in small quantity in the air. 

Hydrogen Dioxide. — Besides water, hydrogen and oxygen 
form a second compound with each other. This is hydro- 
gen dioxide. It is prepared by treating barium dioxicie 



68 INTRODUCTION TO CHEMISTRY. 

with sulphuric acid. The reaction which takes place will 
be explained under barium dioxide. 

Hydrogen dioxide is a liquid which breaks up readily 
into water and oxygen. The ease with which it gives up 
oxygen makes it a good oxidizing agent. It is now manu- 
factured on the large scale. 

Analysis has shown that hydrogen dioxide contains rela- 
tively twice as much oxygen as water does. While, in the 
latter substance, hydrogen and oxygen are combined in the 
proportion of one part by weight of hydrogen to eight parts 
by weight of oxygen, in hydrogen dioxide there are six- 
teen parts by weight of oxygen to one part by weight of 
hydrogen. 

Summary. — We have thus learned that (1) water can be 
decomposed into hydrogen and oxygen by means of an elec- 
tric current; (2) the gases are obtained in the proportion 
of eight parts by weight of oxygen to oue part by weight 
of hydrogen, or one volume of oxygen to two volumes of 
hydrogen; (3) when hydrogen is burned water is formed; 
(4) when hydrogen and oxygen are mixed together they do 
not combine under ordinary circumstances; (5) when a 
spark or flame is brought in contact with the mixture vio- 
lent action takes place accompanied by explosion; (6) the 
action is occasioned by the chemical combination of the 
two gases; (7) they combine in the same proportions as 
those in which they are obtained from water by the action 
of the electric current; (8) water can be made by passing- 
hydrogen over heated copper oxide; (9) by weighing the 
copper oxide before and after the experiment, and deter- 
mining the weight of the water formed, oxygen is found 
to form eight ninths of water. 

Comparison of Hydrogen and Oxygen. — Hydrogen and 
oxygen are different forms of matter, just as heat and 
motion are different forms of energy. Heat can be con- 






COMPARISON OF HYDROGEN AND OXYGEN. 69 

verted into motion, and motion into heat, but one element 
cannot by any means known to us be converted into another. 
They are apparently entirely independent of each other. 
The question will therefore suggest itself, whether, in 
spite of their apparent independence, there is not some re- 
lation between the different elements which reveals itself 
by similarity in properties. It will be found that the ele- 
ments can be divided into groups or families according to 
their properties. There are some elements, for example, 
which in their chemical conduct resemble oxygen markedly. 
These elements make up the oxygen family. So far as hy- 
drogen is concerned, however, it stands by itself. There 
is no other element which conducts itself like it. If we 
compare it with oxygen, we find very few facts which indi- 
cate any analogy between the two elements. In their 
physical properties they are, to be sure, similar. Both are 
transparent, colorless, inodorous gases. On the other hand, 
oxygen combines readily with a large number of sub- 
stances with which hydrogen does not combine. Oxygen, 
as we have seen, combines easily with carbon, sulphur, 
phosphorus, and iron. It is a difficult matter to get 
any of these elements to combine directly with hydrogen. 
Further, substances that combine readily with hydrogen 
do not combine readily with oxygen. The two elements 
exhibit opposite chemical properties. What one can do 
the other cannot do. This oppositeness of properties is 
favorable to combination; for not only do hydrogen and 
oxygen, with their opposite properties, combine with great 
ease under the proper conditions, but, as we shall see later, 
it is a rule that elements of like j>roperties do not readily 
combine with one another, while elements of unlike prop- 
erties do readily combine with one another. 



CHAPTEE V. 

LAWS OF CHEMICAL COMBINATION.— COMBINING 
WEIGHTS.— ATOMIC WEIGHTS.— CHEMICAL EQUATIONS. 

Law of the Indestructibility of Matter. — The earlier 
chemists do not appear to have been very strongly im- 
pressed by the importance of the weight of substances. 
They seem tacitly to have held that matter can be de- 
stroyed or brought into being. The work of Lavoisier, 
however, in the last part of the last century, showed that 
whenever matter is apparently destroyed, it continues to 
exist in some other form. If it were possible to annihilate 
matter or to call it into being at will, it would be of little 
or no scientific value to weigh things. Innumerable experi- 
ments performed since Lavoisier's time have confirmed the 
view that matter is indestructible. The first fundamental 
law bearing upon the changes in composition which the 
different forms of matter undergo is the laiv of the inde- 
structibility of matter, or the laiv of the conservation of 
mass. While it is perhaps impossible to conceive that this 
great law should not be true, it must not be forgotten that 
the only way by which its truth could be established was 
by experiment. The law may be stated thus : 

Whenever a change in the composition of a substance 
takes place the amount of matter after the change is the 
same as before the change. Assuming that this law has 
always held good, it follows that the amount of matter in 
the universe is the same to-day as it has been from the 
beginning. Transformations are constantly taking place, 

70 



LAW OF THE CONSERVATION OF ENERGY. 71 

but these involve no increase nor decrease in the total 
amount of matter. 

Law of the Conservation of Energy. — Just as matter is 
neither created nor destroyed, so it has been shown that 
the total amount of energy is unchangeable. One of the 
greatest discoveries in science is that one form of energy 
can be transformed into others, and that in these trans- 
formations nothing is lost. It is now known that for a 
certain amount of heat a certain amount of motion can be 
obtained, and that for a certain amount of motion a cer- 
tain amount of heat can be obtained. It is known that a 
similar definite relation exists between heat and electrical 
energy. It is known that a definite amount of heat is 
obtained by burning a definite amount of a given substance, 
and it is known also that with a definite amount of heat 
a definite amount of chemical change can be produced. 
Investigation has shown that all the different forms of 
energy are convertible one into the other without loss. 
This great fact is known as the lata of the conservation of 
energy. Transformations of energy are constantly taking 
place, as transformations of matter are, but the total 
amount in each case remains the same. 

Law of Definite Proportions. — Under oxygen the fact 
was mentioned that magnesium and oxygen combine with 
each other in definite proportions. This raises the ques- 
tion as to the proportion by weight in which other ele- 
ments combine with one another. A magnet of a certain 
strength will support a piece of iron of a certain weight. 
But it will also support any piece of iron weighing less. 
It shows no preference for certain weights of iron. So, 
also, the earth attracts all bodies, light or heavy, showing 
no preference for certain weights. When substances act 
upon one another chemically, however, it is found that a 
certain weight of one will combine with a definite weight 



72 INTRODUCTION TO CHEMISTRY. 

of another, and only with this weight — no more and no 
less. Take, for example, the case of iron and sulphur. If 
equal weights of these elements are mixed and caused to 
act chemically by the aid of heat, it will be found that 
some of the sulphur is left in the uncombined state 
after the action is over. If twice as much iron as sulphur 
is taken, then, after the action, some iron is left. If a 
large number of experiments should be made with great 
care, it would be found that when the two elements are 
mixed in the proportion of 7 parts by weight of iron to 
4 parts of sulphur the action is perfect, neither iron nor 
sulphur being left. 

An extensive examination has shown conclusively that 
any given chemical compound always contains the same 
elements in exactly the same proportions. The compound 
of sulphur and iron always contains exactly 36.36 per cent 
of sulphur and 63.64 per cent of iron. The compound of 
magnesium and oxygen always contains exactly 60 per cent 
of magnesium and 40 per cent of oxygen, and so on 
throughout the list of chemical elements. These facts 
were discovered by the united efforts of a large number of 
chemists continued through many years. They are of 
very great importance. They are summed up in the gen- 
eral statement : 

Chemical combination ahv ays talces place between definite 
masses of substances. 

This is known as the law of definite proportions. It is 
simply a statement of what we have every reason to believe 
to be the truth. Every fact known to us in regard to 
chemical combination is in accordance with this general 
statement. It expresses what we learn by a study of chem- 
ical facts. It must be borne in mind that this law, as well 
as other laws governing natural phenomena, can never be 
proved to be absolutely true, for the reason that we cannot 
examine every case to which the law applies. But if, after 
examining a very large number of cases, we find that the 



LAW OF MULTIPLE PROPORTIONS. 73 

law always holds true, we are justified in concluding that 
it is true of all cases. When we say that all bodies attract 
one another, do we know this to be absolutely true ? Cer- 
tainly not. But we do know that, so far as those bodies 
are concerned which come under our observation, the state- 
ment is true, and we therefore have every reason to believe 
that it is true of all bodies. 



Law of Multiple Proportions. — It does not require a very 
extended study of chemical phenomena to show that from 
the same elements it is possible in many cases to get more 
than one product. Thus, iron and sulphur form three 
distinct compounds with each other. Tin combines with 
oxygen in two proportions. The elements potassium, 
chlorine, and oxygen combine in four different ways, form- 
ing four distinct products. Nitrogen and oxygen form 
five products. In the early part of this century the 
English chemist Dalton by a study of cases like those 
mentioned was led to the discovery of another great law of 
chemistry, known as the law of multiple proportions. 
Many substances had been analyzed before his time, and 
the percentage of the constituents determined with a fair 
degree of accuracy. He examined first two gases, both of 
which consist of carbon and hydrogen. He determined 
the percentages of their constituents, and found them to be 
as follows : 

Olefiant gas, 85.7$ carbon and 14.3$ hydrogen; 
Marsh-gas, 75.0$ carbon and 25.0$ hydrogen. 

On comparing these numbers he found that the ratio 
of carbon to hydrogen in olefiant gas is as 6:1; whereas 
in marsh-gas it is as 3 : 1 or 6 : 2. The mass of hydrogen, 
combined with a given mass of carbon, is exactly twice 
as great in the one case as in the other. 



74 INTRODUCTION TO CB3M1STBY. 

There are, further, two compounds of carbon and oxygen, 
and in analyzing these the following figures were obtained: 

Carbon monoxide, 42.86$ carbon and 57.14$ oxygen ; 
Carbon dioxide, 27.27$ carbon and 72.73$ oxygen. 

But 42.86 : 57.14 :: 3:4, and 27.27 : 72.73 :: 3 : 8. 

The mass of oxygen combined with a given mass of 
carbou in carbon dioxide is exactly twice as great as the 
mass of oxygen combined with the same mass of carbon in 
carbon monoxide. These facts and other similar ones led 
to the discovery of the law of multiple proportions, which 
may be stated thus : 

If t zoo elements form several compounds toith each other, 
the masses of one that combine ivitli a fixed mass of the other 
bear a simple ratio to one another. 

The three compounds of iron and sulphur may serve as 
further illustrations. In one of them, approximately 7 
parts by weight of iron are in combination with 4 parts of 
sulphur ; in a second, 7 parts of iron are in combination 
with 6 parts of sulphur ; and in the third, 7 of iron are in 
combination with 8 of sulphur. The figures 4, 6, and 8 
jjlainly bear a simple ratio to one another. The five com- 
pounds of the element nitrogen with oxygen contain 7 
parts of nitrogen and 4, 8, 12, 16, and 20 parts of oxygen 
respectively, which figures plainly bear a simple relation to 
one another, viz., 1:2:3:4:5. 

The law of multiple proportions, like the law of definite 
proportions, is simply a statement in accordance with what 
has been found true by experiment. Although discovered 
by Dalton at the beginning of this century and put for- 
ward upon what appears now to be a slight basis of facts, 
all work since that time has confirmed it, and it forms to- 
day one of the corner-stones of the science of chemistry. 

Combining Weights of the Elements.— A careful study 
of the figures representing the composition of chemical 



COMBINING WEIGHTS OF THE ELEMENTS. 75 

compounds reveals a remarkable fact regarding the relative 
quantities of one and the same element which enter into 
combination with different elements. The proportions by 
weight in which some of the elements combine chemically 
with one another are given in the following table : 



1 part H 


ydrogen combi 


nes 


\vi 


tb 


35.4 


parts Chlorine. 


1 


11 


" ii 




" 




80 


Bromine. 


1 


" 


" 




'« 




126.5 


" Iodine. 


35.4 


parts 


i Chlorine covubi 


ne 


with 


39 


" Potassium. 


80 




Bromine 


" 




" 


39 


• < 


126.5 


" 


Iodine 


" 




" 


39 


" 


16 


" 


Oxygen 


" 




11 


65 


" Zinc. 


16 


" 


" 


< i 




" 


24 


Magnesium. 


16 


<< 


" 


»* 




" 


40 


Calcium. 


16 


« < 


1 1 


< ( 




" 


137 


" Barium. 


65 


it 


Zinc 


" 




« < 


32 


" Sulphur. 


24 


1 1 


Magnesium 


" 




« < 


32 


a it 


40 


< < 


Calcium 


1 1 




** 


32 


i ( n 


137 


< < 


Barium 


" 




1 1 


32 


it tt 



It will be seen that the figures that express the relative 
weights of chlorine, bromine, and iodine that combine with 
1 part of hydrogen also express the relative weights of 
these elements that combine with 39 parts of potassium. 
So also the figures that express the relative weights of zinc, 
magnesium, calcium, and barium that combine with 16 
parts of oxygen also express the relative weights of these 
elements that combine with 32 parts of sulphur. Now 
an examination of all compounds has shown that hydrogen 
enters into combination with the other elements in the 
smallest proportions ; and this element is therefore taken 
as unity in stating the relative weights of the other 
elements that enter into combination. That weight of 
another element that combines with 1 part by weight of 
hydrogen may be called its combining weight. Thus, 
according to this, the combining weights of chlorine, 
bromine, and iodine are respectively 35.4, 80, and 126.5. 
Similarly, 39 is the combining weight of potassium, as it 



76 INTRODUCTION TO CHEMISTRY. 

expresses the weight of potassium that combines with the 
above weights of chlorine, bromine, and iodine. For 
every element a number can be selected such that the 
proportions by weight in which the element enters into 
combination with others can be conveniently expressed by 
this number or by a simple multiple of it. These numbers 
are called the combining weights. 

Hypothesis and Theory. — The laws presented in this 
chapter are then simply condensed statements that sum 
up what has been found true in all cases examined. They 
are statements of facts discovered by experiment. 

When we have established a law by means of experiments, 
the next thing in order is to imagine a cause. We try to 
imagine a condition of things which, if it existed, would 
lead to the results discovered. If we succeed in imagining 
such a condition of things, this leads to an hypothesis. If, 
now, we test this hypothesis in every way that suggests 
itself, and find that all facts discovered are in accordance 
with it, we then call it a theory. An hypothesis is a guess 
in regard to the cause of certain phenomena. A theory is 
an hypothesis which has been thoroughly tested, and 
which is applicable to a large number of related phe- 
nomena. 

Hypotheses and theories are of great value to science, if 
founded upon a thorough knowledge of the facts to which 
they relate. They become dangerous when used by those 
who are not familiar with the facts. Those whose minds 
have not been properly trained are apt to be given to un- 
profitable speculation. The student who has not received 
a thorough scientific training should remember that theories 
and hypotheses, to be of value, must be suggested, not by a 
superficial but by a thorough knowledge of facts. 

With these words of warning and of explanation in re- 
gard to the relation existing between i\\s fact, the law, the 
hypothesis, and the theory, we may proceed to consider 



THE ATOMIC THEORY. 77 

briefly a theory concerning the constitution of matter 
which grew out of the discovery of the law^s of definite 
and multiple proportions. 

The Atomic Theory. — If we consider any simple form of 
matter or element, such as iron, it is clear that there are 
two views which we may hold regarding the way the sub- 
stance is made up. We know we can subdivide every piece 
of iron we can see, no matter how small it may be; and 
though after, a time the particles might become so small 
that we could no longer subdivide them, still we can imagine 
that by more refined methods the process of subdivision 
might be continued without end. If we believe that such 
infinite subdivision is possible, we hold the hypothesis that 
matter is infinitely divisible. We cannot prove this — we 
can only speculate in regard to it. But we may also con- 
ceive that after the process of subdivision has been carried on 
for a time, until very minute particles have been obtained, 
a limit can be reached beyond which the process of sub- 
division cannot be carried. If we believe this, we hold the 
hypothesis that matter is not infinitely divisible, or that 
matter consists of indivisible particles. These particles 
may be called atoms (from the Greek aro/AOS, which sig- 
nifies simply indivisible). Both of these hypotheses have 
been held for ages. But the discussion in regard to the 
relative merits of the two views was at first not much more 
profitable than it would be if carried on between two 
students who are in the early stages of their study of the 
facts. 

AVhen the law 7 s of definite and multiple proportions were 
discovered by Dalton, he saw that the conception that mat- 
ter is made up of indivisible particles or atoms might have 
some connection with the law T s. If each element is made 
up of atoms, the most probable view is that every atom of 
any particular element is exactly like every other atom of 
that element. Among the properties possessed by these 



78 INTRODUCTION TO CHEMISTRY. 

atoms must be weight. It is probable that the atoms of 
different elements have different weights. Suppose now 
that, when chemical combination takes place between two 
elements, the action takes place between these atoms, so 
that one atom of the one element combines with one of the 
other, and soon through the mass. If there were present 
in one mass exactly as many atoms as in the other, both 
substances would be used up — nothing would be left over. 
But if there were a larger number of atoms of one element 
than of the other, then, of the element of which the larger 
number of atoms is present, some would be left over after 
the action is complete. Suppose, further, that the weights 
of the atoms of two elements are to each other as 1 : 10. 
Then, if, when these two elements are brought together, 
they combine in the proportion of one atom of one to one 
atom of the other, the resulting compound would contain 
the elements in the proportion of one part by weight of one 
to ten parts by weight of the other. Or if, on analyzing a 
compound of two elements, we find that it contains one 
part by weight of one to ten parts by weight of the other, 
we should conclude that the weights of the atoms of the 
two elements bear to each other the ratio 1 : 10. 

If matter consists of atoms, and chemical action takes 
place between these atoms, we can understand why chemical 
action takes place between definite weights of substances; 
in other words, we see a probable reason for the law of 
definite proportions. As the atoms are supposed to be in- 
divisible, if two elements combine in more than one pro- 
portion with each other, they must do so in the proportion 
of one atom of one to two atoms of the other, or one to 
three, or two to three, or in some other way which does not 
involve the breaking-up of the atoms. If, for example, two 
elements, the weights of whose atoms are as 1 to 10, com- 
bine in the proportion of one atom of one to one atom of 
the other, the resulting compound will contain the ele- 
ments in the proportion of one part by weight of one to 



ATOMIC WEIGHTS. 79 

ten parts by weight of the other element. If the same ele- 
ments combine in the proportion of one atom of the first to 
two atoms of the other, then the resulting compound will 
contain the elements in the proportion of one jDart by 
weight of one to twenty parts by weight of the other, and 
so on. It will thus be seen that if two elements combine 
in more than one proportion with each other, and the view 
that matter consists of atoms of definite weight, and that 
chemical action takes place between these atoms, is correct, 
then it follows that the elements must combine in accord- 
ance with the law of multiple proportions. 

Atomic Weights. — A. thorough study of the facts has 
shown that the atomic theory, as suggested by Dalton, is 
the simplest conception that can be formed in regard to the 
constitution of matter which will satisfactorily account for 
the laws of definite and multiple proportions. The weights 
of the elements which have thus far been referred to as 
combining weights are, in accordance with the theory, the 
relative weights of the atoms, or the atomic iceights. The 
symbols of the elements represent atoms of the elements. 
Thus H represents an atom of hydrogen, an atom of oxy- 
gen, CI an atom of chlorine, etc. The combining weights, 
found by analyzing compounds in which these elements 
occur, are H = 1, = 16, and CI — 35.4. That is to say, 
by means of these figures we can always represent the rela- 
tive weights of the elements found in their compounds. 
Hydrochloric acid, for example, contains hydrogen and 
chlorine in the proportion of 1 part hydrogen to 35.4 parts 
chlorine. Hence it is believed that the weight of the 
atom of hydrogen is to that of chlorine as 1 to 35.4. As 
hydrogen enters into combination in smaller proportion 
than any other element, its combining weight or atomic 
w T eight is taken as the unit, and all others compared with 
it. If we say that the atomic weight of oxygen is 16, and 
that of chlorine is 35.4, we mean simply that the atom of 



80 INTRODUCTION TO CHEMISTRY. 

oxygen is 16 times heavier and that of chlorine 35.4 times 
heavier than that of hydrogen. We might take any other 
standard, but that of the hydrogen atom is the simplest. 
At one time the atomic weight of oxygen was taken as 100, 
and then the atomic weights of the other elements were 
relatively larger. 

[ Problem. — If we call the atomic weight of oxygen 100, what 
would those of hydrogen and chlorine be ? The atomic weight of 
hydrogen being accepted as 1, those of oxygen and chlorine are 
16 and 35.4 respectively.] 

How the Relative Weights of the Atoms are Deter- 
mined. — If we could isolate atoms and weigh them, there 
would be no serious difficulty in determining their relative 
weights. But as we cannot deal with atoms, we must deal 
with masses of atoms, and from a study of these masses 
draw conclusions regarding the weights of the atoms. 

If it were the rule that two elements combine with 
each other in only one proportion, it might be safe to con- 
clude that they combine in the proportion of one atom of 
one to one atom of the other. Then, by simply determin- 
ing the relative weights of the elements contained in a 
mass of the compound, we should be in a position to draw 
a conclusion regarding the relative weights of the atoms. 
But suppose two elements combine in more than one pro- 
portion. Suppose, for example, that nitrogen and oxygen 
combine, as they do, in these proportions: 14 of nitrogen 
to 8 of oxygen, 7 of nitrogen to 8 of oxygen, 7 of nitrogen 
to 16 of oxygen, and it is required from these figures to 
determine the relative weights of the atoms of nitrogen 
and oxygen. We may suppose that in the first compound 
the elements are combined atom to atom, then the relative 
weights of these atoms are 14 for nitrogen to 8 for oxygen. 
If, however, we had already concluded from a study of 
the compounds of hydrogen and oxygen that the atom of 
oxygen is 16 times heavier than that of hydrogen, we 



FORMULAS OF CHEMICAL COMPOUNDS. 81 

should have in the above compound of nitrogen and 
oxygen 28 parts of nitrogen combined with 16 parts of 
oxygen, and the atomic weight of nitrogen would appear 
to be 28. But we may equally well assume that in this 
compound 2 atoms of nitrogen are combined with 1 atom 
of oxygen. This idea would be represented by the formula 
N 2 0, and, if we accept this conception, the atomic weight 
of nitrogen must be 14. This example will suffice to show 
that the determination of the relative weights of atoms by 
means of the analyses of compounds is a difficult matter, 
and that attempts to make the determinations in this way 
would necessarily lead us into difficulties which we could 
not surmount without the aid of some new conception 
which will aid us in judging of the number of atoms con- 
tained in the smallest particles of compounds. The diffi- 
culties have been largely overcome, as will be shown farther 
on, and the atomic weights accepted at the present day 
have been determined by the aid of a number of methods. 
Much labor has also been expended in making these de- 
terminations, and the problem remains one of the most 
important in chemistry. 

Formulas of Chemical Compounds. — Molecules. — Chem- 
ical compounds are represented by placing the symbols of 
the constituent elements side by side. Thus HC1 means a 
compound of hydrogen and chlorine in which these ele- 
ments are present in the proportion of 1 part by weight of 
hydrogen to 35.4 parts by weight of chlorine, or in terms 
of the atomic theory it means a compound whose smallest 
particle is made up of an atom of hydrogen and an atom 
of chlorine. The formula H 2 stands for a compound 
whose smallest particle is made up of two atoms of hydro- 
gen and one of oxygen; and H 2 2 stands for a compound 
whose smallest particle consists of two atoms of hydrogen 
and two atoms of oxygen. The small figure placed to the 
right below the symbol of an element shows the number of 



82 INTRODUCTION TO CHEMISTRY. 

atoms of the element in the smallest particle of the com- 
pound, and, of course, and this is of chief importance at 
this stage, it shows the proportion by weight in which the 
element is contained in the compound. These smallest 
particles of compounds are called molecules. The relation 
between these and atoms, and the methods of determining 
molecular weights, will be discussed farther on. HC1 
represents then a molecule of hydrochloric acid made up of 
one atom of hydrogen and one atom of chlorine; H a O rep- 
resents a molecule of water; HgO, a molecule of mercuric 
oxide; Mn0 2 , a molecule of manganese dioxide; KC10 3 , a 
molecule of potassium chlorate; ZnS0 4 , a molecule of zinc 
sulphate. These formulas are of great convenience in 
representing chemical reactions. 



CHAPTER VI. 

STUDY OF THE REACTIONS EMPLOYED IN THE PREPA 
RATION OF OXYGEN AND OF HYDROGEN AND IN THE 
STUDY OF WATER, 

Preparation of Oxygen. — The reactions employed in the 
preparation of oxygen were: (1) The decomposition of 
mercuric oxide by heat; (2) The decomposition of potas- 
sium chlorate by heat; (3) The decomposition of man- 
ganese dioxide by heat; and (4) The action of heat on a 
mixture of potassium chlorate and manganese dioxide. 

Heating Mercuric Oxide. — When mercuric oxide is heated 
it is decomposed, yielding mercury and oxygen. That was 
shown in Exp. 3. But analysis shows that mercuric oxide 
consists of mercury and oxygen combined in the propor- 
tion of their atomic weights, 200 parts of mercury to 16 
parts of oxygen; and it has been shown that when decom- 
position takes place, mercury and oxygen are obtained in 
these proportions. These facts are represented by the 
simple equation 

HgO = Hg + 0. 

This equation expresses the reaction qualitatively and 
quantitatively; and it must be remembered that it differs 
from algebraic equations in this important respect, that it 
expresses something that has been established by experi- 
ment. Chemical equations cannot be solved by mental 
processes alone, as algebraic equations can. 

83 



I 

84 INTRODUCTION TO CHEMISTRY. 

Heating Potassium Chlorate. — Potassium chlorate is 
made up as represented by the formula KC10 3 , or its 
molecule consists of an atom of potassium, an atom of 
chlorine, and three atoms of oxygen. The atomic weights 
of these-elements are respectively K = 39; CI = 35.4; and 
= 16. That is to say, the compound is made^of these 
elements in the proportion of 39 parts of potassium; 35.4 
parts of chlorine; and 48 (3 X 16) parts of oxygen. 

It is, therefore, an easy matter to calculate how much 
oxygen, or chlorine, or potassium any given weight of po- 
tassium chlorate contains. Let it be required, for example, 
to calculate how much oxygen is contained in 4 grams of 
potassium chlorate. As the compound is made up of 39 
parts of potassium, combined with 35.4 parts, of chlorine 
and 48 parts of oxygen, in 39 + 35.4 -f 48 = 1^.4 parts of 
potassium chlorate there are 48 parts of oxygen. If in 
122.4 parts there are 48 parts, how much is there in 4 
grams ? Plainly the answer is given by the solution of the 
simple proportion 

122.4 : 48 :: 4 : x, 

in which x represents the actual weight of oxygen con- 
tained in 4 grams of potassium chlorate. Similarly the 
proportion 

122.4 : 39 :: 4 : x 

will give the weight of potassium, and 

122.4 : 35.4 ::4 :x 

will give the weight of chlorine contained in 4 grams of 
potassium chlorate. 

When potassium chlorate is heated until no more oxygen 
is given off, the compound potassium chloride, of the for- 
mula KC1, is left behind. By weighing the potassium chlo- 
rate taken and the potassium chloride left, and measuring 






EXPERIMENTS WITH POTASSIUM CHLORATE. 85 

the oxygen given off, it has been shown that the relative 
weights of the substances are represented by the equation 

KCIO3 = KC1 + 30. 

39 + 35.4 + 48 39 + 35.4 3x16 



122.4 74.4 48 

That is to say, 122.4 grams of potassium chlorate gives 74.4 
grams of potassium chloride and 48 grams of oxygen. 
The figure 3 before the symbol of oxygen means, of course, 
three atoms of oxygen. When the element is in combina- 
tion, the figure expressing the number of atoms is placed 
to the right of the symbol, below the line, as in the formula 
of potassium chlorate. 

Experiment 50. — To determine how much oxygen is given off 
when a known weight of potassium chlorate is decomposed by 
heat, proceed as follows : In a small dry hard-glass tube about 10 
cm. (4 in.) long and 8 to 10 mm. (about \ in.) internal diameter, 
closed at one end, weigh out on a chemical balance about 0.2 
gram dry potassium chlorate, first weighing the tube empty. In- 
troduce just above the potassium chlorate a plug of glass-wool, 
then soften the tube in a flame, and draw it out so that it has 
the form shown in Fig. 26, the plug of glass-wool being at the 

B 




Fig. 26. 
constricted part B of the tube. Now weigh the tube again. 

Let a = weight of tube empty ; 

6 = weight of tube with potassium chlorate ; 

c — weight of tube with potassium chlorate and plug. 

Connect at A by means of a short piece of rubber tubing with 
a measuring-tube (see Fig. 21) so that the ends of the two tubes 
are almost in contact with each other, the measuring-tube having 
been previously filled with water to the zero-point, and the top 



86 INTRODUCTION TO CHEMISTRY. 

closed by means of the stop-cock. Open the stop-cock, and now 
heat the potassium chlorate, gently at first, and gradually higher 
until no more gas is given off. After the gas has stood for half 
an hour to cool it down to the temperature of the air, adjust the 
two tubes of the measuring-apparatus so that the level of the 
water in both is the same ; read off the volume of the gas. At 
the same time read the barometer and thermometer ; and now 
make the corrections for pressure and temperature as directed 
pages 54-57. The weight of a litre or 1000 cc. of oxygen at 0° 
and 760 mm. pressure is 1.4298 grams. Knowing the volume of 
oxygen obtained, calculate the weight of this volume. Remove 
the tube containing the product left after the decomposition of 
the potassium chlorate, and weigh it. 

Let d = weight of tube after decomposition of potassium chlo- 
rate. 

We have 

b — a — weight of potassium chlorate used, 
and 

d — (a + c — 6) = weight of potassium chloride left. 

Knowing further the weight of the oxygen obtained in the de- 
composition, which weight we may call e, it is obvious from what 
has been said that 

d — (a + c — b) + e should be equal to b — a, 

and the weights found should all be in accordance with the equa- 
tion 

KC10 3 = KC1 + 30. 

Make all the calculations, and see how nearly the results ob- 
tained agree with what is represented by this equation. Should 
the results not be satisfactory the first time, repeat the work. 
The more carefully the work is done the more nearly will the 
results agree with the equation. 

In Exp. 17 a fact was observed which is not taken 
account of in the equation KC10 3 = KC1 -f- 30. It was 
seen that the gas was given off in two stages: first, a pare 
came off at a comparatively low temperature, and then a 
larger quantity came off at a higher temperature. If the 
gas given off during the first stage had been measured, it 
would have been found to be only one third of the total 









EXPERIMENTS WITH POTASSIUM CHLORATE. 87 

obtained by complete decomposition. If, further, the solid 
substance left behind in the flask had been properly exam- 
ined, it would have been found to consist of two substances, 
one of which was potassium chloride, KC1, and the other a 
compound that contains more oxygen than the chlorate. 
The latter is potassium perchlorate, KC10 4 . The relative 
quantities of the two substances would also have been 
found to correspond to the weights represented by the for- 
mulas KC1 and KC10 4 , i.e., there would have been found 
39 -f- 35.4 = ?4.4 parts of potassium chloride to 39 + 35.4 
+ 4 x 16 = 138.4 parts of potassium perchlorate. The 
following equation expresses these facts : 

2KC10 3 = KC1 + KC10 4 + 20. 

The figure 2 placed before the formula of potassium 
chlorate affects the whole formula, so that the quantitative 
relations are represented thus : 

2KCIO3 = KC1 + KC10 4 + 20. 

2(3*9 + 35.4 + 48) 39 + 35.4 39 + 35.4 + 64 2x16 

244.8 74.4 138.4 32 

In the second stage of the decomposition all the rest of 
the oxygen is given off, or, in other words, the potassium 
perchlorate is now decomposed, thus : 

KC10 4 = KC1 + 40. 

[Problems. — How much potassium chlorate must be taken to 
get 10 litres of oxygen ? In this case how much potassium per- 
chlorate and how much potassium chloride would be formed ? 
How much potassium chloride would 5 grams of potassium chlo- 
rate yield ? How much potassium perchlorate ? What volume of 
oxygen would be obtained by heating 20 grams of potassium 
chlorate until the first stage of the decomposition is complete ?] 

Heating Manganese Dioxide. — The change effected in 
manganese dioxide by heat is represented by the equation 

3MnO, = Mn 3 4 + 20. 



88 INTRODUCTION TO CHEMISTRY. 

The atomic weight of manganese being 55, the quanti- 
tative relations may be readily calculated. 

[Problem. — How much oxygen can be obtained by heating 12 
grams of manganese dioxide? How much manganese dioxide 
must be heated in order to get 3 grams of oxygen ? In each case 
how much of the compound Mn 3 4 would be obtained ? 

The Action of Oxygen on Carbon, Sulphur, Phosphorus, 
and Iron. — The general character of the action of oxygen 
on the elements named has been discussed in connection 
with Experiments 21, 22, 23, and 24, and evidence was 
presented to show that the action consists in direct union 
which results in the formation of new compounds called 
oxides. But the quantitative relations were not referred 
to. These relations must, of course, be determined by 
experiment. The substance burned, the oxygen used up, 
and the product formed must in each case be weighed. 
This has been done repeatedly, and it will suffice here to 
give the results. 

When sulphur burns in oxygen, for every 32 parts of 
sulphur burned 32 parts of oxygen are used up, and 64 
parts of sulphur dioxide, S0 2 , are formed: 

S + 20 = SO^ 
32 2 X 16 32 + 2 X 16 

32~ ~~ 64 

In the case of carbon it has been shown that for every 12 
parts of carbon burned 32 parts of oxygen disappear, and 
there are formed 44 parts of the compound carbon dioxide, 
C0 2 . The equation representing the action is 

C + 20 = CO,. 
12 2 X 16 12 + 2 X 16 

In the case of phosphorus, for every 62 parts of this ele- 
ment which disappear, 80 parts of oxygen are used up, and 
142 parts of the compound P 2 0. are formed, as represented 
in the equation : 



PREPARATION OF HYDROGEN. 89 

2P + 50 = P 2 5 . 

2 X 31 5 X 16 2 X 31 + 5 X 16 

The compound P 2 6 is known as phosphorus pentoxide. 
It is the white substance found in the vessel. 

When iron burns in oxygen the product formed is mag- 
netic oxide of iron, Fe 3 4 : 

3Fe + 40 = Fe 3 4 . 

3 X 56 4 X 16 3 X 56 + 4 X 16 

168 64 168 64 

Preparation of Hydrogen. — The reactions by which hy- 
drogen can be most readily prepared are: (1) The decom- 
position of water by an electric current; (2) The action of 
sodium and of potassium on water; (3) The action of iron 
on water ; (4) The action of carbon on water ; (5) The 
action of metals on acids. 

Decomposition of Water by an Electric Current. — This 
reaction has been studied somewhat fully in connection 
with the subject of water. The facts established show 
that it is represented by the equation 

H 2 = 2H + 0. 

Action of Sodium and Potassium on Water. — Substitution. 
— The fact that when sodium is thrown on water hydrogen 
is given off was shown in Exp. 27. It was also shown that 
there is something in the water after the action is over. 
Analysis of this substance has shown that it has the 
composition NaOH. It is called caustic soda or sodium 
hydroxide. The reaction between sodium and water is 
represented by the equation 

Na + H 2 = NaOH + H. 

According to this, each molecule of water gives up one 
atom of hydrogen and an atom of sodium takes its place. 
The sodium is substituted for hvdrosren in the water. 




90 INTRODUCTION TO CHEMISTRY. 

This action is plainly different in kind from those which 
have thus far been studied. These are either simple acts 
of combination, as in the action of oxygen on sulphur ; or 
of decomposition, as in the action of heat on mercuric 
oxide and on potassium chlorate. The action of sodium 
on water, however, is an act of substitution, which involves 
both combination and decomposition. The water is de- 
composed and the sodium hydroxide is formed by combi- 
nation. Potassium acts upon water in the same way as 
sodium, and the reaction is represented thus : 

K + H 2 = KOH + H. 

Potassium acts on water more rapidly than sodium does, 
so that the amount of heat evolved in a given time is 
greater than in the case of sodium, and the temperature 
is therefore raised high enough to set fire to the hydrogen. 
If the motion of the sodium is interfered with by placing 
on the water a piece of filter-paper and throwing the metal 
on this, the temperature becomes high enough to set fire 
to the gas, as in the case of potassium. This restriction of 
the motion of the sodium simply prevents it from being 
cooled off, as it is when it moves over the surface of the 
water. 

Action of Iron on Water. — At ordinary temperatures 
iron does not readily act on water, but when steam is 
passed over heated iron, as in Exp. 28, action takes place 
according to this equation: 

3Fe + 4H 2 = Fe 3 4 + 8H. 

(In what other experiment has the compound Fe 3 4 
been obtained ? What is its name ?) 

[Problems. — The atomic weight of iron is 56 ; how much water 
can be decomposed by 20 kilograms (or 40 pounds) of iron? 
and how much would the hydrogen obtained weigh ? One litre 
of hydrogen at 0° C. and under the standard pressure of 760 mm. 
weighs 0.089578 gram ; what will be the volume of hydrogen ob- 



ACTION OF METALS ON ACIDS. 01 

tained by using up 20 kilograms of iron in the decomposition of 
water ?] 

Decomposition of Water by Carbon. — This decomposition 
gives the mixture known as " water-gas." The action is 
represented by the equation 

C + H 2 = CO + 2H. 

The gas CO is carbon monoxide, which will be studied 
later. 

[Problem. — The reaction represented by the last equation 
yields equal volumes of carbon monoxide and of hydrogen. How 
much water would 5 kilograms (or 10 pounds) of carbon decom- 
pose, and what volume of gas would be obtained?] 

Action of Metals on Acids. — The action of zinc on hydro- 
chloric acid and on sulphuric acid was studied in Exp. 29, 
and it was stated that in each case the zinc takes the place 
of the hydrogen in the acid. The two reactions are thus 
represented : 

(1) 2HC1 + Zn = ZnCl 3 + 2H. 

Hydrochloric Acid. Zinc Chloride. 

(2) H 2 S0 4 + Zn = ZuSO, + 2H. 

Sulphuric Acid. Zinc Sulphate. 

In Exp. 30 the zinc sulphate formed was obtained in 
the form of crystals. It will be seen that the action of 
zinc on hydrochloric acid and on sulphuric acid is of 
the same kind as the action of sodium and of potassium 
on water. It is substitution. The zinc takes the place of 
the hydrogen. Another point of interest to be noted is 
that while an atom of sodium takes the place of one atom 
of hydrogen, an atom of zinc takes the place of two atoms 
of hydrogen. It will be seen later that there are elements 
whose atoms have the power of taking the place of three 
atoms of hydrogen, and others with still higher substitut- 
ing- values. 



92 



INTRODUCTION TO CHEMISTRY. 



Experiment 51.— The amount of hydrogen evolved when a 
known weight of zinc is dissolved in sulphuric acid can be de- 
termined by means of the apparatus represented in Fig. 27. The 




Fig. 27. 

bent tube leading from the flask A is drawn out at B, and a plug 
of glass-wool introduced below the constriction. The other parts 
of the apparatus need no description. The flask should have a 
capacity of about 40 to 50 cc. ; and the measuring-tube C should 
have a capacity of about 100 cc, and be graduated in T \ cc. Fill 
D with distilled water that has been boiled ; put a piece of zinc 
weighing from 0.15 to 0.20 gram in the flask ; open the pinch-cock E, 
by which means the whole apparatus is filled with water. Examine 
the apparatus to see whether gas-bubbles are lodged under the 
stopper i^or in the glass-wool. If so, they can usually be dislodged 
without difficulty. If they persist, boil the water for a few mo- 
ments. Now place the measuring-tube C in position and let the 
greater part of the water remaining in D flow through the ap- 
paratus. Into this tube D then pour sulphuric acid (1 of acid 
to 4 of water) until it is nearly full. Open the pinch-cock E, and 
thus displace the water which fills the apparatus. The action of 
acid upon the metal is facilitated by heat or by adding with the 
zinc a few small pieces of platinum. When the action is over, 
sweep the contents of the flask through the delivery-tube by again 
opening the pinch-cock E. Finally, the measuring-tube is trans- 
ferred to a cylinder of water (see Fig. 22), and the volume of the 



PREPARATION OF HYDROGEN DIOXIDE. 93 

gas read and corrected in the usual manner. How much does the 
hydrogen obtained in the experiment weigh ? How much ought 
to have been obtained ? How many cubic centimetres ought to 
have been obtained ? 

Action of Hydrogen on Copper Oxide. — In Exp. 46 it was 
shown that when hydrogen is passed over heated copper 
oxide, water is formed and copper is left in the tube. The 
action is represented by the equation 

CuO + 2H = H a O + Cn. 

Here, it will be observed, two atoms of hydrogen take the 
place of one atom of copper in the oxide. (Compare the 
action of zinc on sulphuric acid with that of hydrogen on 
copper oxide.) 

[Problem. — The atomic weight of copper is 63.2; how much 
water would be formed by reducing 5 grams of copper oxide? 
How much hydrogen would be necessary?] 

Preparation of Hydrogen Dioxide.— Double Decomposition. 

— Hydrogen dioxide consists of hydrogen and oxygen com- 
bined in the proportion of 1 part by weight of hydrogen to 
16 parts by weight of oxygen. By methods which will be 
discussed later it will be shown that the molecule of this 
compound consists of two atoms of hydrogen combined with 
two atoms of oxygen, its formula being H 2 2 . Its forma- 
tion from barium dioxide by the action of sulphuric acid is 
represented thus: 

Ba0 2 + H 2 S0 4 = Ba-S0 4 + H 2 2 . 

Barium Dioxide. Sulphuric Acid. Barium Sulphate. 

In some respects this reaction differs from all others thus 
far studied. Here, two compounds acting upon each other 
give two new compounds. There is an exchange of constit- 
uents. This kind of action is called double decomposition, 
or metathesis. It is by far the most common kind of chem- 
ical action with which we have to deal, 



94 INTRODUCTION TO CHEMISTRY. 

Kinds of Chemical Reactions. — There are then four kinds 
of chemical reactions, and all tfiese have been illustrated by 
examples. They are: (1) Direct Combination; (2) Direct 
Decomposition; (3) Substitution; and (4) Double Decompo- 
sition, or Metathesis. 

Conditions under which Chemical Reactions Take Place. — 

Chemical reactions take place under the greatest variety of 
conditions. Some take place by simply bringing the sub- 
stances together at ordinary temperature, as, for example, 
when sodium and potassium act on water. Others take place 
on heating the substances together, as, for example, when 
iron decomposes water, and hydrogen decomposes copper 
oxide. In most of the reactions thus far studied, and in- 
deed in most of those which will be studied, heat is em- 
ployed, as it generally aids chemical action. Some reactions 
take place in solution. That this is a great aid to chemical 
action has already been pointed out (see page 66). Some 
substances are so unstable that they are decomposed with 
violence by the touch of a feather, as, for example, a com- 
pound of iodine and nitrogen. Others are so stable that 
they resist the action of the highest temperatures. The 
electric current has a very marked effect upon many chem- 
ical compounds, especially if they are in solution or in 
molten condition. This has been illustrated by the action 
of the current on water, the result being, as will be remem- 
bered, the setting free of the two gases oxygen and hydro- 
gen. Again, reactions differ very much in respect to their 
violence, from the most terrific explosions to the quiet ac- 
tion of the air in our lungs, and of carbon dioxide in the 
leaves of plants. 



CHAPTER VII. 

CHLORINE AND ITS COMPOUNDS WITH HYDROGEN 
AND OXYGEN. 

Occurrence. — Chlorine, though widely distributed in 
nature, does not occur in very large quantity as compared 
with oxygen and hydrogen. It is found chiefly in combi- 
nation with the element sodium as common salt, or sodium 
chloride, which has the composition represented by the 
formula NaCl. It is also found in combination with other 
elements, as potassium, magnesium, etc. In comparatively 
small quantity it occurs in combination with silver, form- 
ing one of the most valuable silver ores. All the chlorine 
with which we have to deal is made from common salt. 

Preparation. — It is not practicable to decompose sodium 
chloride directly into its elements. In order to get the 
chlorine out of the compound in the free state, it is neces- 
sary, first, to get it in combination with hydrogen in the 
form of hydrochloric acid, HC1. This is very easily ac- 
complished by treating salt with ordinary sulphuric acid. 
When the two are brought together a change takes place, 
which will be studied more in detail farther on. The re- 
action is represented thus : 

(1) 2NaCl + H 2 S0 4 = Na 2 S0 4 + 2HC1. 

As will be seen, the sodium of the sodium chloride and the 
hydrogen of the hydrochloric acid exchange places, a kind 

95 



96 INTRODUCTION TO CHEMISTRY. 

of action which is quite common. This particular reaction 
is of very great importance in the arts, as it is the first 
stage in the preparation of common "soda" or sodium 
carbonate, and of chlorine. 

Now, if hydrochloric acid is brought in contact with a 
substance which gives up oxygen easily, the hydrogen unites 
with the oxygen to form water, and the chlorine is set 
free. The reaction is expressed thus : 

(2) 2HC1 + = H 2 + 2C1. 

[Problem. — How much sulphuric acid will be required to set 
free enough hydrochloric acid to make 25 grams of chlorine ?] 

Deacon's Process. — As there is an unlimited supply of 
oxygen in the air, it would be advantageous to effect the 
decomposition of hydrochloric acid by means of the ele- 
ment in the free state. On the large scale this is now 
accomplished. Deacon's process for manufacturing chlo- 
rine consists in passing air and hydrochloric acid together 
through a heated tube containing clay balls saturated with 
copper sulphate. Exactly why the oxidation takes place 
under these circumstances is not known. The essential 
feature of the reaction is nevertheless the oxidation of the 
hydrochloric acid, as represented in equation (2). 

Laboratory Method. — For the preparation of chlorine in 
the laboratory it is most convenient to bring hydrochloric 
acid in contact with manganese dioxide, Mn0 2 , a substance 
which has been employed for the purpose of preparing oxy- 
gen. The action which takes place is explained thus: In 
the first place, when hydrochloric acid acts upon many 
compounds containing oxygen, the hydrogen and oxygen 
combine, and the element which was in combination with 
oxygen combines with chlorine. Thus, when the com- 
pound MnO is treated with hydrochloric acid, this reac- 
tion takes place : 

MnO + 2HC1 = MnCl 2 + H 2 0. 



PREPARATION OF CHLORINE. 97 

So, also, when manganese dioxide is treated with hydro- 
chloric acid, the oxygen is probably first replaced by chlo- 
rine, as represented in the equation 

Mn0 2 + 4HC1 = MnCl 4 + 2H 2 0. 

But the compound MnCl 4 gives up half of its chlorine 
when heated: 

MnCl 4 = MnCl 2 + 2C1. 

So that the action of hydrochloric acid on manganese di- 
oxide is represented as follows : 

Mn0 2 + 4HC1 = MnCl 2 + 2H 2 + 2C1. 

[Problem. — How much manganese dioxide would be required 
to liberate 50 grams of chlorine ? The combining weight of man- 
ganese is 55.] 

Instead of first making the hydrochloric acid from salt 
and then treating the hydrochloric acid with manganese 
dioxide, it is simpler to mix together the manganese di- 
oxide and common salt and pour upon the mixture the 
necessary quantity of sulphuric acid. The reaction takes 
place according to the following equation: 

Mn0 2 + 2NaCl + 2H 2 S0 4 

= Na 2 S0 4 + MnS0 4 + 2H 2 + 2C1. 

The products formed are sodium sulphate, Na 2 S0 4 , man- 
ganese sulphate, MnS0 4 , water, H 2 0, and chlorine. 

When sulphuric acid acts upon sodium chloride alone 
the products are sodium sulphate, Na 2 S0 4 , and hydrochlo- 
ric acid, HC1: 

2NaCl + H 2 S0 4 = Na 2 S0 4 + 2HC1. 

Experiment 52.— Pour 2 or 3 cc. concentrated sulphuric acid 
on a gram or two of common salt in a test-tube. Heat gently. 
Describe all that you observe. Hydrochloric acid is formed. 
Give some account of it. 



98 INTRODUCTION TO CHEMISTRY. 

When sulphuric acid acts upon manganese dioxide alone 
the products are manganese sulphate, MnS0 4 , water, and 
oxygen : 

Mn0 2 + H 2 S0 4 = MnS0 4 + H 2 + 0. 

Experiment 53. — In a small flask provided with a delivery- 
tube heat strongly a few grams (4 or 5) of manganese dioxide 
with enough concentrated sulphuric acid to cover it completely. 
Collect the gas given off, and examine it. Is it oxygen ? 

When manganese dioxide and sodium chloride are heated 
together with sulphuric acid it is possible that both the 
above reactions take place, and that the products hydro- 
chloric acid and oxygen then act upon each other giving 
water. It is, however, also possible that the hydrochloric 
acid acts upon the manganese dioxide as represented in the 
equation on the last page. 

Experiment 54. — Arrange an apparatus as shown in Fig. 28. 
The flask need not, of course, be of the shape shown in the figure 
but may be of any kind available. It should have a capacity of 
about 1 litre. The delivery-tube should extend to the bottom of 
the collecting vessel, which should be a clean, dry cylinder or 
bottle of colorless glass. The mouth of this vessel should be 
covered with a piece of paper to prevent currents of air from 
carrying away the chlorine. The experimenter can by the color 
judge of the quantity of chlorine present in the vessel. A satis- 
factory method of making chlorine is this: Mix 5 parts by weight 
of coarsely granulated common salt and 5 of coarsely granulated 
manganese dioxide. Make a mixture of 12 parts by weight of 
concentrated sulphuric acid and 6 of water.* Let this mixture 
cool down to the ordinary temperature, and then pour it on the 
mixture of manganese dioxide and common salt. Use 25 grams 
manganese dioxide and the other substances in the proper pro- 
portion. Collect six or eight dry cylinders or bottles full of 
chlorine. 

(1) Introduce into one of the vessels containing chlorine a little 
finely-powdered antimony. The two elements combine at once 

* See precautions noted on page 41. 



EXPERIMENTS WITH CHLORINE. 



99 



with evolution of light. The product is antimony trichloride, 
SbCl 3 . 

[In what respects does this experiment resemble the one in 
which iron was burned in oxygen ? In what respects do the two 
differ?] 

(2) Into a second vessel introduce a few r pieces of heated thin 




Fig. 28. 

copper foil. Combination takes place with evolution of light and 
heat. 

(3) Into a third vessel introduce a piece of paper with some 
writing on it, some flowers, and pieces of colored calico. Most 
of the colors will be destroyed if the substances are moist. 

(4) Into a fourth vessel introduce a dry piece of the same 
colored calico as that used in the previous experiment. What 
difference do you observe between the dry piece and the moist 
piece ? 



100 INTRODUCTION TO CHEMISTRY. 

Properties of Chlorine. — Chlorine is a greenish-yellow 
gas. It has a disagreeable smell, and acts upon the pas- 
sages of the throat and nose, causing irritation and inflam- 
mation. The effect is much like that of a " cold in the 
head." Inhaled in concentrated condition, i.e., not diluted 
with a great deal of air, it would cause death. It is much 
heavier than air, its specific gravity being 2.45. A litre of 
chlorine gas, under standard conditions, weighs 3.167 
grams. It is soluble in water, acts upon mercury, and 
therefore cannot be collected by displacement of either of 
these liquids. The most convenient way to collect it is by 
displacement of air as in the experiment. 

Action of Chlorine. — From these experiments it is seen 
that chlorine combines readily with other substances, and 
also that it destroys colors, or bleaches. It is indeed one 
of the most active elements. It not only acts directly 
upon many of the elements at ordinary temperatures, and 
decomposes many compounds, but it also acts upon most 
organic substances, or such as are formed as the products 
of animal or vegetable life. Its action upon the tissues of 
the respiratory organs has already been noticed. 

Experiment 55. — Cut a piece of filter-paper about an inch 
wide and six to eight inches long. Pour on this some ordinary 
oil of turpentine previously warmed slightly. Introduce this 
into a vessel of chlorine. What takes place ? 

Oil of turpentine consists of carbon and hydrogen. The 
main action of the chlorine in this case consists in extract- 
ing the hydrogen and leaving the carbon. The experi- 
ment is interesting chiefly in so far as it illustrates the 
general tendency of chlorine to act upon vegetable sub- 
stances. 

Bleaching by Chlorine. — It has been noticed that when 
moisture is present chlorine bleaches, while when it is 
not present bleaching does not take place, It has been 



CHLORINE HYDRATE. 



101 



3 



shown that chlorine acts directly upon some dye-stuffs, 
converting them into colorless substances. In other cases 
it has been shown that the destruction of the color is 
due to oxygen, which is set free from water by the action 
of chlorine. In the direct sunlight chlorine 
decomposes water according to this equation : 

2C1+H 2 = 2HCl + 0. 

In bleaching, this decomposition of water 
takes place in direct contact with the colored 
materials, and the oxygen, the instant it is 
set free, is more active than free oxygen. It 
is this oxygen which is being set free that 
acts upon the colored substances and con- 
verts them into colorless substances. 

Experiment 56. — Seal the end of a glass tube 
about a metre (or about a yard) long and about 
12 mm. (i inch) internal diameter. Fill this with 
a strong solution of chlorine in water. Invert it, Fig. 29. 
as shown in Fig. 29, in a shallow vessel containing some of the 
same solution of chlorine in water. Place the tube in direct 
sunlight. What takes place ? What change in color of the solu- 
tion, in the odor, taste ? Examine the gas. What does it appear 
to be? 




Chlorine Hydrate. — When chlorine gas is passed into 
water cooled down almost to the freezing-point, crystals 
appear in the vessel. These consist of chlorine and water 
and are known as chlorine hydrate. Its composition is 
represented by the formula CI + 5H 2 0. The crystals are 
very unstable, breaking up at the ordinary temperature 
into chlorine ?as and water. 



Experiment 57. — Light a jet of hydrogen in the air and care- 
fully introduce it into a vessel containing chlorine. Does it con- 
tinue to burn ? Is there any change in the appearance of the 
flame? Is there any evidence of the formation of a gas that 



102 INTRODUCTION TO CHEMISTRY. 

differs from hydrogen and from chlorine ? In what other experi- 
ments has this gas been met with ? 

Hydrogen Burns in Chlorine. — The ease with which 
chlorine decomposes water and vegetable substances con- 
taining hydrogen shows that it readily combines with hy- 
drogen. Just as hydrogen burns in oxygen, it also burns 
in chlorine. 

The burning of hydrogen in air or oxygen is simply the 
act of combination of hydrogen and oxygen, the product 
being water in the state of vapor, and therefore invisible. 
When hydrogen burns in chlorine the action consists in 
the union of the two gases, the product being hydrochloric 
acid, HC1, which forms the clouds in the air. In both 
cases the action is accompanied by an evolution of heat 
and light. 

Chlorides. — Just as the compounds of oxygen with other 
elements are called oxides, so the compounds of chlorine 
with other elements are called chlorides. 

Nomenclature of Chlorides and of Oxides. — The chlorides 
are named in the same way as the oxides. The name of 
the element with which the chlorine is combined is pre- 
fixed. Thus the compound of zinc and chlorine, ZnCl 2 , is 
called zinc chloride; that of sodium and chlorine, sodium 
chloride, etc. When an element forms more than one 
compound with oxygen, suffixes are made use of to dis- 
tinguish between them. Thus in the case of copper there 
are two oxides which have the formulas Cu 2 and CuO. 
The former, which contains the smaller proportion of oxy- 
gen, is called cuprous oxide, while the latter, which contains 
the larger proportion of oxygen, is called cupric oxide. In 
general, of two oxides of the same element, that which 
contains the smaller proportion of oxygen is designated by 
the suffix ous, while that which contains the larger pro- 
portion is designated by the suffix ic. Ferrous oxide has the 



RELATION OF LIGHT TO CHEMICAL ACTION. 103 

composition FeO; ferric oxide } Fe 2 3 ; manganous oxide is 
MnO; manganic oxide is Mn 2 3 . In case there are more 
than two oxides the number of atoms of oxygen in the 
molecule of the compound is frequently indicated in the 
name. Thus manganese dioxide is Mn0 2 ; sulphur tri- 
oxide is S0 3 ; phosphorus pent oxide is P 2 5 , etc. Chlorides 
are named in the same way. Cuprous chloride is CuCl; 
cupric chloride is CuCl 2 ; ferrous chloride is FeCl 2 ; ferric 
chloride is FeCl 3 ; phosphorus trichloride is PC1 3 ; selenium 
tetrachloride is SeCl 4 , etc. 

Hydrochloric Acid, HC1. — The only compound that chlo- 
rine and hydrogen form with each other is hydrochloric 
acid. It has already been shown that hydrogen burns in 
chlorine and that hydrochloric acid is formed. The two 
gases may be mixed and allowed to stand together in- 
definitely in the dark and no action will take place. If, 
however, the mixture is put in diffused sunlight, gradual 
combination takes place; and if the direct light of the sun 
is allowed to shine for an instant on the mixture, explosion 
occurs, and this is the sign of the combination of the two 
gases. The same sudden combination is effected by apply- 
ing a flame or spark to the mixture, or by illuminating it 
instantaneously with an electric light or the light from a 
piece of burning magnesium. 

[What difference is there between the combination of 
hydrogen and oxygen and of hydrogen and chlorine ?] 

Relation of Light to Chemical Action. — The way in 

which the sunlight and other bright lights act upon the 
mixture of hydrogen and chlorine to cause them to com- 
bine is not understood; but the fact that sunlight does 
have a marked influence upon some kinds of chemical 
action is well known. One other illustration of this in- 
fluence has already been before us, that of the decom- 
position of water by chlorine. This action does not take 



104 INTRODUCTION TO CHEMISTRY. 

place in the dark. The sunlight is essential. The art of 
photography is based upon the influence of light in causing 
chemical changes. The light from the object photo- 
graphed is allowed to act in the camera on a plate, upon 
the surface of which is a substance which is changed chemi- 
cally by light. It should be specially noted that the cause 
of the chemical changes in these cases is not the heat, but 
the light. If the substances are heated to the same tem- 
perature in the dark, the changes do not take place. 

Preparation of Hydrochloric Acid. — To prepare hydro- 
chloric acid, common salt or sodium chloride, NaCl, is 
treated with sulphuric acid (see Experiment 52, p. 97). 
As has already been explained, the hydrogen of the sul- 
phuric acid and the sodium of the salt exchange places, as 
represented in the equation 

2NaCl + H 2 S0 4 = Na 2 S0 4 + 2HC1. 

The products are sodium sulphate and hydrochloric acid. 
The hydrochloric acid is given off as a gas, and the sodium 
sulphate remains behind in the flask. 

Properties. — Hydrochloric acid is a colorless transparent 
gas. It has a sharp penetrating taste and smell. If in- 
haled into the lungs it produces suffocation. It dissolves 
in water very readily. At ordinary temperatures one volume 
of water will dissolve 450 times its own volume of the gas. 
The solution is the liquid known in the laboratory as hy- 
drochloric acid. 

[Problem. — A litre of hydrochloric acid gas weighs 1.6283 grams 
at 0°. At 0° one volume of water will absorb 500 times its own 
volume of the gas. How much will a litre of water increase in 
weight at 0° by taking up all the hydrochloric acid it can ?] 

So readily does hydrochloric acid combine with water 
that it condenses moisture from the air; hence, although 
the gas itself is quite colorless and transparent, when it 



HYDROCHLORIC ACID. 105 

comes in contact with the air dense white clouds are 
formed, which are not formed if it is kept from contact 
with the air. — Hydrochloric acid does not burn and does 
not support combustion. This is equivalent to saying that 
it does not combine with oxygen under ordinary circum- 
stances, and that substances which combine with the oxy- 
gen of the air do not combine with hydrochloric acid. 

[What evidence have you had that, under some circum- 
stances, oxygen does act on hydrochloric acid ? What are 
the circumstances ? What are the products ?] 

Commercial hydrochloric acid is a yellowish liquid, the 
color being due to the presence of impurities. The solu- 
tion is obtained in the factories in which " soda," or sodium 
carbonate, is made. This is an extremely important sub- 
stance in the arts. It does not occur in nature, but is 
manufactured from common salt. In the process com- 
monly used the salt is first converted into sodium sul- 
phate by treating it with sulphuric acid. Hydrochloric 
acid is necessarily given off. When the factories were first 
established in England, the gas was allowed to escape as a 
waste-product, but the effects produced by it upon the 
vegetation of the surrounding countrv were so destructive 
that a law was passed prohibiting the manufacturers from 
allowing the gas to escape. It is now collected by passing 
it through water. Thus enormous quantities of the solu- 
tion are produced, but its uses are numerous and it always 
commands a price. 

Pure hydrochloric acid is a solution of the pure gas in 
pure water. It is colorless, and when concentrated it 
gives off fumes when exposed to the air. The solution 
when heated gives off a large part of the gas contained in 
it, and by boiling it can all be evaporated. 

Experiment 58. — Arrange an apparatus as shown in Fig. 30. 
The flask may be any ordinary one of about 1 litre capacity. 



106 



INTRODUCTION TO CHEMISTRY. 



The tubes leading into the Wolff's bottles must not dip in the 
water in the bottles. If they end a few millimetres above the 
surface of the water all the gas will be absorbed. 

Weigh out 5 parts common salt, 5 parts concentrated sulphuric 
acid, and 1 part water. Mix the acid and water, taking the 
usual precautions ; let the mixture cool down to the ordinary 
temperature ; and then pour it on the salt in the flask. For the 
purposes of the experiment take about 25 grams of salt. Now 
heat the flask gently, and the gas will be regularly evolved. 
What does the fact of the sinking of the solution through the 
water indicate ? — After the gas has passed for ten to fifteen min- 
utes, disconnect at A, and let some escape into the air. Explain 
what you see. Blow your breath towards the end of the tube. 
Does this produce any effect? Explain this. Apply a lighted 
match to the end of the tube. Does the gas burn ? — Collect some 
of the gas in a dry cylinder by displacement of air, as in the 
case of chlorine. The specific gravity of the gas being 1.26, the 
vessel must of course be placed with the mouth upward. That 




Fig. 30. 

the gas is colorless and transparent is shown by the appearance 
of the generating-flask, which is filled with the gas. Insert a 
burning stick or candle in the cylinder filled with the gas. What 
takes place ? Does the gas support combustion ? Reconnect the 
generating-flask with the series of bottles containing water, and 
let the process continue until no more gas comes over. The re- 
action represented in the equation 

2NaCl 4- H 2 S0 4 = Na 2 S0 4 + 2HC1 



HYDROCHLORIC ACID. 107 

is now complete. After the flask has cooled down, pour water 
on the contents ; and when the substance is dissolved filter it 
and evaporate to such a concentration that, on cooling, the 
sodium sulphate is deposited. Pour off the liquid and dry the 
solid substance by placing it upon folds of filter-paper. Compare 
the substance with the common salt which you put in the flask 
before the experiment. What proofs have you that the two sub- 
stances are not the same ? — Heat a small piece of each in a dry 
tube closed at one end. What differences do you notice ? — Treat 
a small piece of each in a test-tube with sulphuric acid. What 
difference do you notice? — If in the experiment you should re- 
cover all the sodium sulphate formed, how much would you 
have? Put about 50 cc. of the liquid from the first W^olfTs 
bottle in a porcelain evaporating-dish. Heat over a small flame 
just to boiling. Is hydrochloric acid given off?— Can all the 
liquid be driven off by boiling? — Try the action of the solution 
on some iron filings. What is given off ? — Add some to a little 
granulated zinc in a test-tube. What is given off ? Add a little 
to some manganese dioxide in a test-tube. What is given off ? — 
Add ten or twelve drops of the acid to 2 to 3 cc. water in a test- 
tube. Taste the dilute solution. How would you describe the 
taste ? — Add a drop or two of a solution of blue litmus, or put 
into it a piece of paper colored blue with litmus. The color is 
changed to red. Litmus is a vegetable color prepared for use as 
a dye. Other vegetable colors are changed by hydrochloric acid. 
— Steep a few leaves of red cabbage in water. Add a few drops 
of the solution thus obtained to dilute hydrochloric acid. Is there 
any change in color ? — In each case add a few drops of a solution 
of caustic soda. What change takes place ? — In what experiment 
has caustic soda been obtained ? What relation does it bear to 
water ? — To the dilute solution of hydrochloric acid add drop by 
drop a dilute solution of caustic soda. What change takes place 
in the taste ? 

Analysis of Hydrochloric Acid. — The determination of 
the composition of hydrochloric acid is not as easily made 
as that of water. That it consists of hydrogen and chlo- 
rine is shown by the fact that it is formed by direct com- 
bination of these elements. To determine the relative 
weights and volumes of the gases which enter into com- 



108 INTRODUCTION TO CHEMISTRY. 

bination, we may proceed thus : Enclose a suitable quan- 
tity of the gas in a tube. Introduce a small piece of the 
metal potassium. Decomposition will take place as rep- 
resented in the equation 

K + HOI = KOI + H. 

The gas left over is hydrogen. On measuring its volume 
it will be found to be just half that of the hydrochloric 
acid decomposed. The weight of the hydrogen obtained 
will be found to bear to the weight of the hydrochloric 
acid the proportion 1 : 36.4. In other words, in 36.4 parts 
of hydrochloric acid there are 35.4 parts of chlorine and 1 
part of hydrogen. In 1 volume of the gas there is -J vol- 
ume of hydrogen. By mixing equal volumes of hydrogen 
and chlorine and causing them to combine it has been 
found that 1 volume of hydrogen combines with 1 volume 
of chlorine to form 2 volumes of hydrochloric acid. The 
specific gravity or the relative weights of equal volumes of 
hydrogen and chlorine are: hydrogen, 0.0691; chlorine, 
2.45. These figures bear to each other the same relation 
as the atomic weights of the elements, viz., 1 : 35.4. [What 
fact of the same kind was noticed on comparing the spe- 
cific gravities of hydrogen and oxygen ?] Regarding the 
chemical conduct of hydrochloric acid, the experiments 
already performed have shown: 

1. That it gives up its hydrogen when brought in contact 
with certain substances like iron, zinc, etc., which belong 
to the class called metals; and that it takes up the metals 
in place of the hydrogen. Thus zinc and hydrochloric 
acid give zinc chloride and hydrogen : 

Zn + 2HC1 = ZnCl 2 + 2H. 

2. That in contact with substances which give off 
oxygen, or with oxygen itself under certain circumstances, 
it gives up its chlorine, while the hydrogen combines with 
oxygen to form water. 



COMPOUNDS OF CHLORINE. 109 

It will be seen hereafter that when it acts upon the com- 
pounds of the metals with oxygen or the so-called metallic 
oxides like magnesia or magnesium oxide, MgO; lime or 
calcium oxide, CaO ; zinc oxide, ZnO, etc., — compounds 
which do not easily give up oxygen, — the hydrogen of the 
acid combines with the oxygen of the oxide to form water, 
while the metals combine with the chlorine : 

MgO + 2HC1 = MgCl 2 + H 3 0; 
CaO + 2HC1 = CaCl 2 + H 2 0; 
ZnO + 2HC1 = ZnCl 2 + H 2 0. 

It will be noticed that when hydrochloric acid acts upon 
zinc oxide, zinc chloride is formed. But this is the product 
obtained when hydrochloric acid acts upon the metal zinc. 
The metals calcium and magnesium act towards hydro- 
chloric acid the same as zinc. Plainly the cause of these 
reactions is the tendency on the part of chlorine to unite 
with the metallic elements. 

Compounds of Chlorine with Oxygen and with Hydrogen 
and Oxygen. — As has been seen, chlorine combines very 
readily with hydrogen, and hydrogen with oxygen, and the 
products are stable compounds. On the other hand, chlo- 
rine cannot be made to combine directly with oxygen. By 
indirect processes they can be combined, but the com- 
pounds undergo decomposition easily, yielding back the 
chlorine and oxygen contained in them. Before taking 
up these compounds it will be well to study, as far as may 
be necessary, the compounds of chlorine, hydrogen, and 
oxygen which are more easily made, and from which the 
oxides are made. 

Compounds of Chlorine with Hydrogen and Oxygen. — 

One of the principal reactions made use of for the prepa- 
ration of compounds of chlorine, oxygen, and hydrogen is 
that which takes place when potassium hydroxide is 
treated with chlorine. It has been shown that chlorine 



110 INTRODUCTION TO CHEMISTRY. 

combines readily with metals. Now, when chlorine is 
brought together with potassium hydroxide we should ex- 
pect it to combine with the potassium thus: 

KOH + CI = KC1 + + H. 

But its tendency to combine with hydrogen would cause 
the two to unite, so that the result would, in the first stage, 
be represented thus : 

KOH + 2C1 = KC1 + HOI + 0. 

The oxygen, however, can combine with potassium chlo- 
ride, KC1, to form compounds K010, KC10 2 , KC10 3 , and 
KC10 4 ; and the hydrochloric acid formed would combine 
with potassium hydroxide thus: 

KOH + HC1 = KC1 + H 2 0. 

By treating potassium hydroxide with chlorine we may 
therefore expect to obtain potassium chloride, KC1; some 
compound of potassium chloride with oxygen; an' water. 
The above equations are given with the view of making 
clear what actually takes place, as has been shown by exper- 
iment. The products are different according to circum- 
stances. If the solution of caustic potash is warm and 
concentrated, the reaction takes place as represented thus : 

6K0H + 601 = 5KC1 + KC10 3 + 3H 2 0. 

A part of the potassium chloride is oxidized to the form 
KC10 3 , which is known as potassium chlorate. 

If, however, the solution is dilute, the reaction takes 
place thus: 

2K0H + 201 = KC1 + KC10 +H 2 0. 

In the latter case the oxidation of the potassium chloride 
is not carried as far as in the former. The product KC10 
is known as potassium hypochlorite. 

Potassium chlorate, KC10 3 , and potassium hypochlorite, 
KC10, bear the same relation to two compounds, H010 9 



COMPOUNDS OF CHLORINE. Ill 

and HC10, that potassium chloride, KC1, bears to hydro- 
chloric acid, HC1, or sodium chloride, NaCl, to hydrochlo- 
ric acid. But hydrochloric acid can be very easily obtained 
from sodium chloride by simply adding sulphuric acid, 
and potassium chloride undergoes the same change when 
treated with sulphuric acid. Further, it will be shown 
later that nearly all compounds containing sodium or 
potassium give up these elements when treated with sul- 
phuric acid, and take up hydrogen in their place. 

On treating potassium chloride with sulphuric acid this 
reaction takes place : 

2KC1 + H 2 S0 4 = K 2 S0 4 + 2HC1. 

Similarly, on treating potassium chlorate with sulphuric 
acid, this reaction takes place: 

2KC10 3 + H 2 S0 4 = K 2 S0 4 + 2HCIO3. 

The products are potassium sulphate and chloric acid, 
H010 3 . The chloric acid, however, is very unstable, and 
decomposes, yielding other compounds of chlorine. The 
acid itself can be made by taking proper precautions, but 
the chief interest connected with it is the fact that it de- 
composes very easily. Potassium chlorate, which is so 
closely related to it, is a very important compound. As 
has been shown, it gives up its oxygen under the influence 
of heat. It also gives up oxygen in contact with many 
substances which have the power to unit with this element. 
It is a powerful oxidizing agent. 

Potassium hypochlorite, KCIO, formed by passing chlo- 
rine into a dilute solution of caustic potash, is decomposed 
by sulphuric acid thus : 

2KC10 + H 2 S0 4 = K 2 S0 4 + 2HC10. 

The products are potassium sulphate and hypochlorous 
acid. If a concentrated solution of potassium hypochlorite 
is treated with sulphuric acid, the hypochlorous acid formed 



112 INTRODUCTION TO CHEMISTRY. 

at once undergoes decomposition, yielding chlorine, water, 
and oxygen. The acid itself is not well known. The prin- 
cipal compound related to it is " bleaching-powder," or the 
substance generally known as " chloride of lime/' which is 
familiar to every one on account of its application as a 
disinfecting agent. This is made by passing chlorine into 
slaked lime, which from a chemical point of view is anal- 
ogous to caustic potash. Just as when chlorine acts on a 
dilute solution of caustic potash a mixture of potassium 
chloride and potassium hypochlorite is formed, so when 
chlorine acts on slaked lime a mixture of calcium chloride, 
CaCl 2 , and calcium hypochlorite, Ca(OCl) 2 , is formed. 
This mixture is bleaching-powder. By treating it with an 
acid it gives up chlorine, and hence it affords a convenient 
means of transporting chlorine. Many thousands of tons 
of this powder are manufactured annually. The chlorine 
is passed into the lime. It is held chemically combined 
until it is wanted, when it can be liberated by adding an 
acid or by exposure to the air. 

Experiment 59. — Dissolve 40 grams (or about 1£ ounces) 
caustic potash in 100 cc. water in a beaker-glass, and pass chlo- 
rine into it. Arrange an inverted funnel on the end of the de- 
livery-tube so that the edge of the funnel dips just below the 
surface of the potash solution. This is to prevent the choking of 
the delivery-tube by the formation of chlorate in it. When chlo- 
rine passes freely through the solution, thus indicating that it is 
no longer absorbed, stop the action. Filter the solution, and 
allow it to cool, when crystals of potassium chlorate will be de- 
posited, mixed with a little potassium chloride. Recrystallize 
from a little water. Filter off the crystals and dry them. What 
evidence have you that the substance is potassium chlorate ? 
Does it give off oxygen when heated ? In a dry test-tube pour 
two or three drops of concentrated sulphuric acid on a small 
crystal of the substance. Do the same with a piece of potassium 
chlorate from the laboratory bottle. Hold the mouth of the test- 
tube away from the face. What is noticed in each case ? — Evap- 
orate the solution from which the crystals of potassium chlorate 



COMPOUNDS OF CHLORINE. 



113 



have been removed. On allowing it to cool crystals will again 
be deposited. Take them out and recrystallize them. Does this 
substance give off oxygen when heated ? Does it give off a gas 
when treated with sulphuric acid? Is this gas colored? Is it 
hydrochloric acid ? How do you know that it is ? If the gas is 
hydrochloric acid, what is the solid substance from w r hich it is 
formed ? And what is left in the test-tube ? 

Experiment 60. — Mix 20 to 30 grams (about 1 ounce) of fresh 
quick-lime w T ith 50 cc. water. After the slaking is over, pass 
chlorine into it until the gas is no longer absorbed, using the 
same form of delivery-tube as in the last experiment, Put the 
powder thus formed in a flask arranged as shown in Fig. 31. 




Fig. 31, 



Pour a mixture of equal parts of sulphuric acid and water slowly 
through the funnel-tube. Collect by displacement of air the gas 
given off. What evidence have you that the gas is chlorine ? 

Decomposition of Bleaching-powder by Acids.— In the last 
experiment the substance first formed is bleaching-pow^der, 
or " chloride of lime." This is decomposed by sulphuric 
acid, yielding chlorine. The formation of chlorine is sec- 
ondary, and due to the ease with which hypochlorous acid 



114 INTRODUCTION TO CHEMISTRY. 

breaks up into chlorine, oxygen, and water. The tendency 
of sulphuric acid to extract calcium, just as it does potas- 
sium, and to put hydrogen in its place, is at the root of the 
matter. Potassium hypochlorite and potassium chloride, 
when treated with sulphuric acid, yield primarily hypochlo- 
rous acid and hydrochloric acid : 

2KC10 + H 2 S0 4 = K 2 S0 4 + 2HC10; 
2KC1 + H 2 S0 4 = K 2 S0 4 + 2HC1. 

Thus far the only change that has taken place is the ex- 
change of hydrogen for potassium. Now, however, the 
hypochlorous acid decomposes, yielding oxygen, water, and 
chlorine, probably thus : 

2H010 = 201 + H 2 + 0. 

The oxygen thus liberated would, however, act upon 
hydrochloric acid, if present, and set chlorine free: 

2HCl + = H a O + 2Cl; 

so that, if a mixture of potassium hypochlorite and potas- 
sium chloride is treated with sulphuric acid, we should 
expect the result to be that which is represented in this 
equation : 

KC10 + KC1 + H 2 S0 4 = K 2 S0 4 + H 2 + 2C1. 

This in reality expresses what takes place, as has been 
proved experimentally. The decomposition of "bleach- 
ing-powder " takes place in the same way, the only differ- 
ence being that in one case we have to deal with compounds 
of the metal potassium, while in the other we have to deal 
with analogous compounds of the metal calcium. 

Other Compounds of Chlorine, Hydrogen, and Oxygen. — 

While the remaining compounds of chlorine, hydrogen, 
and oxygen cannot be considered here in detail, a reference 
to the series as a whole will serve to call to mind some im- 



COMPOUNDS OF CHLORINE. 115 

portant matters of general interest. There are four of 
these compounds which, as far as composition is concerned, 
bear a very simple relation to one another. They are 
hypochlorous acid, HCIO; chlorous acid, HC10 2 ; chloric 
acid, HC10 3 ; and perchloric acid, HC10 4 . Beginning with 
hydrochloric acid, we have thus a series of compounds, the 
successive members of which differ by one atom of oxygen: 

Hydrochloric acid HC1 

Hypochlorous acid HCIO 

Chlorous acid HCIO, 

Chloric acid HC10 3 

Perchloric acid HC10 4 

This series illustrates very clearly the laic of maltijjlc 
proportions (see ante, p. 73). [What is the law of mul- 
tiple proportions ? How does this series illustrate the 
law?] 

Compounds of Chlorine and Oxygen, — There are two of 
these compounds, viz., chlorine monoxide, C1 2 0, and chlo- 
rine dioxide, C10 2 . They are unstable substances which 
easily break up into chlorine and oxygen. They are not 
easily prepared in pure condition. 



CHAPTER VIII. 

ACIDS.— BASES.— NEUTRALIZATION.— SALTS. 

Neutralization. — It is now time to inquire what features 
acids have in common which lead chemists to give them 
that name. It is not possible to understand the nature of 
their common properties without a somewhat premature 
reference to a class of substances to which special attention 
will be called in due time. These are the alkalies, which 
are the most marked representatives of the class of sub- 
stances known as lases. These two classes, the acids and 
the lases, have the power to destroy the characteristic 
properties of each other. When an acid is brought in 
contact with a base in proper proportions, the character- 
istic properties of both the acid and the base are destroyed. 
They are said to neutralize each other. This act of neu- 
tralization is an extremely important one, with which we 
constantly have to deal in chemical operations. 

Litmus Test for Acids and Alkalies. — The most common 
acids are sulphuric, hydrochloric, and nitric acids. Among 
the more common bases are caustic soda, caustic potash, 
and lime. A convenient way to recognize whether a sub- 
stance has acid or basic properties is by means of certain 
color-changes. The dye litmus is blue. If a solution that 
is colored blue with litmus is treated with a drop or two 
of an acid, the color is changed to red. If now the red 
solution is treated with a few drops of a solution of a base, 
the blue color is restored. There are many other sub- 
stances which have markedly different colors in acid and 

116 



EXPERIMENTS ON NEUTRALIZATION. 



117 



in alkaline solutions. An infusion of red cabbage, for 
example, changes color when treated with an acid, and re- 
covers its color when again treated with an alkali. 



! 




Experiment 61. — Make dilute solutions of nitric, hydrochloric, 
and sulphuric acids (4 cc. dilute acid, such as is used in the 
laboratory, to 200 cc. water); and of caustic soda and caustic 
potash (about 1 gram to 200 cc. of water). Measure off about 20 
cc. of each of the acid solutions. Add a few drops of a solution of 
blue litmus. Gradually add to each 
of the measured quantities of acid 
sufficient dilute caustic soda to cause 
the red color just to change to blue. 
As long as the solution is red it is 
acid. When it turns blue it is alka- 
line. At the turning-point it is neu- 
tral. The operation is best carried 
on by means of a burette, which is 
a graduated tube with an opening 
from which small quantities can be 
poured. A convenient shape is that 
represented in Fig. 32. At the 
lower end is a small opening. The 
flow of the liquid from the burette 
is controlled by means of a small 
pinch-cock. It will require some 
practice to enable the student to 
know exactly when the red color 
disappears and the blue appears, but 
with practice the point can be recog- 
nized with great accuracy. Should 
too much alkali be allowed to get 
into the acid, add a small measured 
quantity of the acid from another 
burette. Having in one experiment determined how much of the 
solution of alkali is required to cause the red color to change to 
blue in operating on a given quantity of the acid solution, try 
the experiment again, using a different quantity of the acid solu- 
tion. Perform similar experiments with the other acids. After- 
wards carefully examine the numerical results. 



Fig. 32. 



118 INTRODUCTION TO CHEMISTRY. 

Ratio of Acid to Alkali in Neutralization. — If the results 
of several experiments with the same acid and alkali are 
recorded it will be found that there is a definite ratio be- 
tween the quantities of acid and alkali solution required 
to neutralize one another. If, for example, 15 cc. of the 
alkali solution are required to neutralize 20 cc. of the acid 
solution, 18 cc. of the alkali solution will be required to 
neutralize 24 cc. of the acid solution, 30 cc. to neu- 
tralize 40 cc, etc. In other words, in order to neutral- 
ize a given quantity of an acid, a definite quantity of an 
alkali is necessary. Suppose 15 cc. of the caustic-soda 
solution or 12 cc. of the caustic-potash solution should be 
required to neutralize 20 cc. of the hydrochloric-acid solu- 
tion. If the quantities of these alkali solutions necessary 
to neutralize equal quantities of the other acids are com- 
pared, it will be found that, if it requires 15 cc. caustic- 
soda solution or 12 cc. caustic-potash solution to neutralize 
20 cc. hydrochloric-acid solution, then the quantities of 
caustic-soda solution and caustic-potash solution required 
to neutralize any definite quantity of a solution of another 
acid will be to each other as 15 to 12. 

What is Formed when Acid and Base are Neutralized ? — 

It appears, therefore, from these experiments that the act 
of neutralization is a definite one, which takes place be- 
tween definite quantities of acid and base. The next 
question that suggests itself is, What is formed when the 
acid and base are neutralized ? Experiment must answer. 

Experiment 62. — Dissolve 10 grams caustic soda in 100 cc. 
water. Add hydrochloric acid slowly, examining the solution 
from time to time by means of a piece of paper colored blue with 
litmus. As long as the solution is alkaline it will cause no 
change in the color of the paper. The instant it passes the point 
of neutralization it changes the color of the paper to red ; when 
exactly neutral it will neither change the blue to red, nor, if the 
color is changed to red by means of another acid, will it change 
it back again. When this point is reached, evaporate off the 



NEUTRALIZATION REACTIONS. 119 

water on the water-bath to compl-te dryness, and see what is 
left. Taste the substance. Has it an acid taste ? Does it sug- 
gest any familiar substance ? If it is sodium chloride, how ought 
it to conduct itself when treated with sulphuric acid ? Does it 
conduct itself in this way ? Satisfactory evidence can be given 
that the substance is sodium chloride. It is not an acid nor an 
alkali. It is neutral. Its formation took place according to the 
equation 

HC1 + NaOH = NaCl + H 2 0. 

Using nitric acid and caustic soda, the product formed is 
sodium nitrate. Compare it with sodium nitrate from the labo- 
ratory bottle. Heat a small specimen of each in a tube closed at 
one end. What do you obser*3 ? Treat a small specimen of each 
with a little sulphuric acid in test-tubes. What do you observe ? 

The Reactions. — The explanation of the changes which 
occur in these cases will be given later. Here the point to 
be noted is, that the substance formed when nitric acid 
acts on caustic soda is sodium nitrate. The reaction took 
place thus: 

HN0 3 + NaOH = NaN0 3 + H 2 0. 

Similarly sulphuric acid and caustic soda give sodium 
sulphate and water, thus : 

H 2 S0 4 + 2NaOH = Na,S0 4 + 2H 2 0. 

With caustic potash similar reactions take place. Hy- 
drochloric acid and caustic potash yield potassium chloride 
and water: 

HC1 + KOH = KC1 + H 2 0. 

Nitric acid and caustic potash yield potassium nitrate 
and water : 

HNOs + KOH = KNOs + H 2 0. 

Sulphuric acid and caustic potash yield potassium sul- 
phate and water : 

H 2 S0 4 + 2KOH •= K 2 S0 4 + 2H 2 0. 



ISO INTRODUCTION TO CHEMISTRY. 

What these Experiments Show. — They show: 

(1) That an acid contains hydrogen; 

(2) That a base contains a so-called metallic element; 

(3) That when an acid acts on a base the hydrogen and 
metallic element exchange places; 

(4) That the substance obtained from the acid by replac- 
ing the hydrogen by a metallic element is neutral; 

(5) That the substance formed from the base by substi- 
tuting hydrogen for the metallic element is water. 

These statements are of general application, except state- 
ment (4), to which there are some exceptions. It is true 
in some cases that after replacing the hydrogen the sub- 
stance has an alkaline reaction; and in other cases that 
the product has an acid reaction. 

It has been shown that hydrochloric acid and sulphuric 
acid act upon certain metals, as iron and zinc, and that 
the action consists in giving up hydrogen and taking up 
metal in its place. The products of this action are the 
same in character as those formed by the action of acids 
on bases. 

Definitions of Acids, Bases, and Salts. — An acid is a sub- 
stance containing hydrogen, which it easily exchanges for 
a metal when treated with a metal itself, or with a com- 
pound of a metal, called a base. 

A base is a substance containing a metal combined with 
hydrogen and oxygen. It easily exchanges its metal for 
hydrogen when treated with an acid. 

The products of the action of an acid on a base are, first, 
water, and, second, a neutral substance called a salt. 

In the examples already given sodium chloride, potassium 
chloride, sodium nitrate, potassium nitrate, sodium sul- 
phate, and potassium sulphate are salts. 

Metallic Elements.— It may fairly be asked, What is a 

« metallic element ? Unfortunately for our present purpose, 

it is by no means an easy matter to give a satisfactory 



NOMENCLATURE OF ACIDS. 121 

answer to this question. Examples of metals can easily 
be given, such as iron, zinc, silver, calcium, magnesium, 
etc.; but when the attempt is made to state what the dis- 
tinguishing features of these substances are, difficulties are 
met with. In general, it may be said that to the chemist 
any element is metallic which with hydrogen and oxygen 
forms a product that has the power to neutralize acids; 
that is to say, that has basic properties. In general, any 
element that has the power to enter into an acid in the 
place of the hydrogen is called a metal, or is said to have 
metallic properties. This is the sense in which the word 
metal is used in this book. 

Nomenclature of Acids. — The names of those acids of 
chlorine which contain oxygen illustrate some of the prin- 
ciples of nomenclature in use in chemistry. That acid of 
the series which is best known is called chloric acid. The 
termination ic is generally used in naming acids, as is seen in 
the names hydrochloric, sulphuric, nitric, etc. If a second 
acid containing the same elements exists and the proportion 
of oxygen contained in it is smaller than in the acid the 
name of which ends in ic, the second acid is given a name 
ending in ous. Thus chlorous acid contains a smaller pro- 
portion of oxygen than chloric acid, and the suffixes ic and 
ous signify that fact. There are many other examples of 
this use of these suffixes in the names of acids as well as in 
the names of compounds of other classes. 

In the series of chlorine acids, however, this simple prin- 
ciple, which is sufficient for most cases, does not suffice. 
In order, therefore, to form characteristic names for the 
other members of the series recourse is had to prefixes. 
There is one acid which, so far as the proportion of oxygen 
contained in it is concerned, stands below chlorous acid. 
It is called /^;j>ochlorous acid, the prefix hypo being de- 
rived from the Greek vno, under. Further, there is an 
acid which contains a larger proportion of oxygen than 



i22 INTRODUCTION TO CHEMISTRY. 

chloric acid. It is called jperchloric acid, the Latin prefix 
per signifying here very or fully. It will be seen that the 
names of the acids vary with the proportion of oxygen con- 
tained in them. 

Nomenclature of Bases. — As pointed out above, a base is 
a compound of a metal with hydrogen and oxygen. Thus, 
caustic soda has the formula NaOH, caustic potash KOH, 
lime Ca0 2 H 2 , etc. They are commonly known as hydrox- 
ides. In order to distinguish between the hydroxides of 
the different metals, the names of the metals are put before 
the name hydroxide. Thus, caustic soda, NaOH, is called 
sodium hydroxide; caustic potash, KOH, is called potas- 
sium hydroxide; caustic lime, Ca0. 2 H 2 , is called calcium 
hydroxide, etc. They are regarded as water in which a 
part of the hydrogen has been replaced by a metal, and 
indeed many of them can be made by simply bringing the 
corresponding metals in contact with water. Thus, as has 
been seen (Exp. 27, page 37), when sodium or potassium 
is thrown on water hydrogen is evolved. The products 
formed are, respectively, sodium hydroxide and potassium 
hydroxide. These compounds are called hydrates by some 
chemists, the name implying that they are derivatives of 
water. The name hydroxide means simply that the sub- 
stances contain hydrogen and oxygen. 

Nomenclature of Salts. — Theoretically every metal can 
yield a salt with every acid. The salts derived from a 
given acid receive a general name, and this general name 
is qualified in each case by the name of the metal con- 
tained in the salt. Thus all the salts derived from nitric 
acid are called nitrates; all the salts derived from chloric 
acid are called chlorates; the salts of sulphuric acid are 
called sulphates.* So too, further, the salts of chlorous 

* If the principle were strictly applied the salts of sulphuric acid 
would be called sulphurates, but for the sake of convenience the 
name is shortened. 



NOMENCLATURE OF SALTS. 123 

acid are called clilorites; those of nitrous acid, nitrites; 
those of sulphurous acid, sulphites, etc., etc. It will be 
noticed that the terminal syllable of the name of the salt 
differs according to the name of the acid. If the name of 
the acid ends in ic, the name of the salt derived from it 
ends in ate. If the name of the acid ends in ous, the 
name of the salt ends in ite. To distinguish between the 
different salts of the same acid, the name of the metal con- 
tained in it is prefixed. Thus, the potassium salt of nitric 
acid is called potassium nitrate, the sodium salt is called 
sodium nitrate; the calcium salt of sulphuric acid is called 
calcium sulphate; the magnesium salt of nitrous acid is 
magnesium nitrite. The calcium salt of hypochlorous acid 
is calcium hypochlorite, etc., etc. [Give the name and 
formula of the potassium salt of perchloric acid. — Give the 
name and formula of the sodium salt of hypochlorous acid. 
— Give the name and formula of the sodium salt of chlorous 
acid.] 

If the salts of hydrochloric acid were named in accord- 
ance with the principle just explained, they would be 
called hydrochlorates. But it will be observed that these 
salts are identical with the products formed by direct 
combination of the metals with chlorine. Thus, hydro- 
chloric acid and zinc act as represented in the equation 

Zn + 2HC1 = ZnCl 2 + 2H, 

while zinc and chlorine act thus : 

Zn + 2C1 = ZnCl 2 . 

In each case the same product, ZnCl 2 , is formed. But 
these compounds of metals with chlorine are called chlo- 
rides, as has already been explained. Hence the name 
hydrochlorate is unnecessary. 

Acid Properties and Oxygen. — The observation that 
oxygen is generally present in acids led at one time to the 



1 24 INTROD UGTION TO CHEMISTR Y. 

belief that it is an essential constituent of these substances. 
Hence the name oxygen was given to it (from o£v$, acid, 
and yevvaoQ, I form). That oxygen is not essential to 
the existence of acid properties is shown in the case of 
hydrochloric acid, and in a few other similar cases. It 
must be said, however, that the acid properties of sub- 
stances are generally due to the presence of oxygen. Some 
substances with basic properties can be converted into 
acids by causing them to combine with oxygen. 



CHAPTER IX. 



NITROGEN— AIR. 

Two Gases in the Air. — It has been stated that when 
substances burn in the air the same products are formed as 
when they burn in oxygen; and, further, that there is 
something besides oxygen present in the air which renders 
the burning less active than it is in oxygen alone. When 
phosphorus is exposed to the air it combines slowly with 
oxygen, and the phosphorus pentoxide formed easily dis- 
solves in water. These facts may be utilized for the pur- 
pose of getting possession of the other gas in the air, as 
this does not combine with phosphorus. 

Experiment 63. — Arrange an apparatus as in Fig. 33. Use 
a tube graduated in cubic centimetres. Enclose 
60 to 80 cc. air in the tube over water. Arrange 
the tube so that the level of the water inside and 
outside is the same. Note the temperature 
of the air and the height of the barometer. 
Reduce the observed volume to standard con- 
ditions. Now introduce a piece of phosphorus 
fastened to the end of the wire, and allow it 
to stand for twenty-four hours. Draw out 
the phosphorus. Again arrange the tube so 
that the level of the water inside is the same 
as that outside. Make the necessary correc- 
tions for temperature, pressure, and tension 
of water-vapor. It will be found that the 
volume has diminished considerably, but that 
about four fifths of the gas originally put in 
the tube is still there. If the work is done 
carefully, the volume of the gas left in the FlG - 33 - 

tube will be to the total volume used as 79 to 100. In Othel- 
lo 




126 INTRODUCTION TO CHEMISTRY. 

words, of every 100 cc. air used 21 cc. are absorbed by phos- 
phorus, and 79 cc. are not. The gas absorbed is oxygen, identical 
with the oxygen made from the oxide of mercury, manganese 
dioxide, and potassium chlorate. The gas left over has no 
chemical properties in common with oxygen. Carefully take the 
tube out of the vessel of water, closing its mouth with the thumb 
or some suitable object to prevent the contents from escaping. 
Turn it with the mouth upward, and introduce into it a burning 
stick. What takes place ? This residual gas will not support 
combustion, and cannot therefore be oxygen. 

Nitrogen. — The experiment just performed shows us 
that the air is made up by volume of 21 per cent of oxygen 
and 79 per cent of a gas which does not support combus- 
tion. This second constituent of the air is nitrogen. 

Preparation. — Anything that has the power to absorb 
oxygen may be used in the preparation of nitrogen from 
the air. To avoid contamination of the nitrogen with 
other substances, however, it is necessary to use something 
which does not form a gaseous product when burned. 
Metallic copper is convenient, and is not unfrequently 
used. It is only necessary to pass air over heated copper, 
when the metal combines with oxygen, forming the solid 
copper oxide, CuO, leaving the nitrogen uncombined. 
The most convenient way to prepare nitrogen is to burn a 
piece of phosphorus in a bell-jar over water. 

Experiment 64. — Place a good-sized stoppered bell- jar over 
water in a pneumatic trough. In the middle of a flat cork about 
three inches in diameter fasten a small porcelain crucible, and 
float this on the water in the trough. Put into it a piece of phos- 
phorus about twice the size of a pea, and set fire to it. Quickly 
place the bell-jar over it. At first some air will be driven out of 
the jar. [Why ?] The burning will continue for a short time, and 
then gradually grow less and less active, finally stopping. On 
cooling, it will be found that the volume of gas is less than four 
fifths the original volume, for the reason that some of the air 
was driven out qf the vessel $t the beginning of the experiment, 



THE AIR. 127 

Before removing the stopper of the bell-jar see that the level of 
the liquid outside is the same as that inside. Try the effect of 
introducing successively several burning bodies into the nitrogen, 
— as, for example, a candle, a piece of sulphur, phosphorus, etc. 
If convenient, place a live mouse in a trap in a bell-jar over 
water. When the oxygen is used up the mouse will die. After 
the animal gives plain signs of discomfort, it may be revived by 
taking away the bell-jar and giving it a free supply of fresh air. 

The Air. — The nitrogen and oxygen which make up the 
air are not chemically combined with each other, but 
simply mixed together. It is not an easy matter to prove 
this statement, but the evidence is so strong that no 
chemist doubts it. 

(1) If nitrogen and oxygen are mixed together, the 
mixture conducts itself in exactly the same way as air. 
The mixing is not accompanied by any phenomena indi- 
cating chemical action. Generally, the chemical union of 
two substances is accompanied by a change in their tem- 
perature. When nitrogen and oxygen are mixed there is 
no change in the temperature of the gases. 

(2) The composition of a chemical compound is con- 
stant. The law of definite proportions is founded upon a 
very large number of observations, and in all cases in which 
there is independent evidence that chemical action takes 
place, it is found that the same substances combine in the 
same proportions to form the same product. Variation in 
the composition of a chemical compound is not known. 
The composition of the air varies slightly, according to cir- 
cumstances. 

(3) Air dissolves somewhat in water. If air which has 
been thus dissolved is pumped out and analyzed, it is 
found to have a composition different from that of ordinary 
air. Instead of containing 1 volume of oxygen to 4 vol- 
umes of nitrogen, it contains 1 volume of oxygen to 1.87 
volumes of nitrogen. The relative quantity of the oxygen 
is much larger in air tha,t has been dissolved in water th*vn 



128 INTRODUCTION TO CHEMISTRY. 

it is in ordinary air. This is due to the fact that oxygen 
is more soluble in water than nitrogen is. In order, how- 
ever, that one gas may dissolve more than the other, it is 
necessary that they should not be in chemical combina- 
tion. If they were in chemical combination the compound 
as such would probably dissolve. 

Occurrence of Nitrogen. — Besides being found in the free 
state in the air, nitrogen is found in combination in a large 
number of substances in nature. It is found in the ni- 
trates, or salts of nitric acid, particularly as the potassium 
salt, KN0 3 , known as saltpetre, and the sodium salt, NaN0 3 , 
known as Chili saltpetre. It is also found in the form of 
ammonia, which is a compound of nitrogen and hydrogen, 
represented by the formula NH 3 . Ammonia occurs in 
small quantity in the air, and is formed under a variety of 
conditions, to which reference will be made when the sub- 
stance is presented. Nitrogen occurs, further, in most 
animal substances in chemical combination. 

Properties of Nitrogen. — It has been seen that nitrogen is 
a colorless, tasteless, inodorous gas. It does not support 
combustion, nor does it burn. (Suppose nitrogen were 
combustible, what would be the composition of the atmos- 
phere ?) Nitrogen not only does not combine directly 
with oxygen readily, but it does not combine directly with 
any other element except at very high temperature. Just 
as it does not support combustion, so also it does not 
support respiration. An animal would die in it, not on 
account of any active poisonous properties possessed by it, 
but for lack of oxygen. In the air it serves the useful 
purpose of diluting the oxygen. If the air consisted only 
of oxygen, all processes of combustion would certainly be 
much more active than they now are. What the effect on 
animals of the continued breathing of oxygen would be, it 
is difficult to say. 



OTHER CONSTITUENTS OF THE AIR. 



129 



Other Constituents of the Air. — Besides nitrogen and 
oxygen the air contains other substances, some of which 
are of great importance. 

Experiment 65. — On a watch-glass expose a few pieces of cal- 
cium chloride to the air. What change takes place, and how is 
this explained ? See Experiment 42. (What is a salt called which 
has the power to take up water from the air and dissolve in the 
water ?) 

Experiment 66. — Lime-water is made by putting a few pieces 
of quick-lime in a bottle and pouring water upon it. The mix- 
ture is well shaken and allowed to stand. The undissolved 
substance settles to the bottom, and with care a clear liquid can 
be poured off the top. This is lime-water, which is a solution of 
calcium hydroxide, Ca(OH) 2 , in water. Baryta-water is a solu- 
tion of a similar compound of the metal barium. Expose some 
clear lime-water or baryta- water and note the changes. 

When these solutions are exposed to nitrogen or oxygen, 
or to an artificially prepared mixture of the two gases, no 




Fig. 34. 



change takes place. Further, if air is first passed through 
a solution of caustic soda it no longer has the power to 



130 INTRODUCTION TO CHEMISTRY. 

cause the formation of a crust on lime-water or baryta- 
water. 

Experiment 67. — Arrange an apparatus as shown in Fig. 34. 
The wash-cylinders A and B are half filled with ordinary caustic- 
soda solution. The bottle C is filled with water. The tube D 
reaches to the bottom of the bottle. Being filled with water and 
provided with a pinch-cock, it acts as a siphon. Open the pinch- 
cock and let the water flow slowly out of the bottle. As it flows 
out, air will be drawn in through the caustic-soda solution in the 
wash-cylinders. When the bottle is filled with air pour some 
water in again so that it is about a quarter full. Draw this 
water off as before. Now remove the stopper from the bottle, 
pour in 20 to 30 cc. lime-water and cork the bottle. Is there any 
difference between the action of this air and ordinary air ? What 
difference ? 

What Causes the Difference? — It appears, therefore, that 
there is something present in the air under ordinary cir- 
cumstances which has the power to form a crust on lime- 
water or baryta-water, and which can be removed by pass- 
ing the air through caustic soda. Thorough examination 
has shown that this is the compound called carbon dioxide 
or carbonic acid gas. It is the substance obtained by burn- 
ing charcoal in oxygen. 

Experiment 68. — Into the bottle containing the air from which 
the carbon dioxide has been removed, insert a burning stick or 
taper for a moment. Notice whether a crust is now formed on 
the lime-water. Wood and the material from which the taper is 
made contain carbon. Explain the formation of the crust on the 
lime-water after the stick of wood or taper has burned for a 
short time in the vessel. 

Experiment 69. — Arrange an apparatus as shown in Fig. 35. 
The bottle A contains air. B contains concentrated sulphuric 
acid ; C is carefully dried and contains a few pieces of granulated 
calcium chloride and air. Pour water through the funnel-tube 
into A, the air will be forced through B and into C. But in 
passing through B the moisture contained in it will be removed, 
ftnd the air which, enters C will be dry. After A has once been 



WATER- VAPOR IN THE AIR. 



131 



filled with water, empty it and fill it again, letting the dried air 
pass into C. This operation may be repeated as many times as 




Fig. 35. 

may seem desirable. The calcium chloride in C will not grow 
moist. 

Constituents of the Air. — The preceding experiments 
show that besides oxygen and nitrogen there are present in 
the air water, in the form of vapor, and carbon dioxide, 
which is a colorless gas. Wherever we examine the air 
these two substances are found to be present. Indeed, it 
is evident that they must be present. Evaporation is tak- 
ing place everywhere, even at low temperatures, and the 
vapor thus formed is carried to all parts of the earth 
by the winds. Whenever any of our ordinary combustible 
substances burn in the air, carbon dioxide is formed; and, 
further, the process of respiration of animals also gives rise 
to the formation of carbon dioxide, which is given off from 
the lungs. 

Quantity of Water-vapor in the Air. — The quantity of 
water-vapor present in the air varies between comparatively 
wide limits. At any given temperature the air cannot hold 



132 INTRODUCTION TO CHEMISTRY. 

more than a certain quantity. When it contains this quan- 
tity it is said to be saturated. And if cooled down below 
this temperature the vapor partly condenses and appears 
now as water. When a vessel containing ice-water is placed 
in the air, that which immediately surrounds the vessel is 
cooled down below the point at which the quantity of 
water-vapor present would saturate the air, and water con- 
denses on the outside of the vessel. Every one has noticed 
that on a warm cloudy day more water condenses on a cold 
object than on a clear cool day. The water-vapor present 
in the air has an important effect on man. The inhabi- 
tants of countries with moist climates apparently have 
characteristics which are not generally met with in those 
who inhabit countries with dry climates. The difference 
in the effects of moist and of dry air on an individual is 
well known. 

When air which is charged with water-vapor comes in 
contact with cooler air, the vapor condenses and falls as 
rain. 

The quantity of w T ater-vapor in a given volume of air 
can be determined by drawing the air through a weighed 
tube containing calcium chloride. This will absorb the 
water and increase in weight, and the increase in weight 
will represent the quantity of water in the volume of air 
drawn through the tube. 

Quantity of Carbon Dioxide in the Air. — The quantity 
of carbon dioxide in the air is relatively very small, being 
about 3 parts in 10,000 parts of air. It varies slightly ac- 
cording to the locality and the season. It is essential to the 
growth of plants. 

Other Substances in the Air. — Besides nitrogen and oxy- 
gen, carbon dioxide and water, the air contains a small 
quantity of ammonia (see p. 134), and a large number of 
other substances in very small quantities, 



CHAPTER X. 

COMPOUNDS OF NITROGEN WITH HYDROGEN AND 
OXYGEN. 

General Conditions which Give Rise to the Formation of 
the Simpler Compounds of Nitrogen. — It has been shown 
that nitrogen is an inactive element, manifesting little ten- 
dency to combine with other elements. It is nevertheless 
an easy matter to get compounds of nitrogen with many 
other elements, and among these compounds, some of those 
w T hich it forms with hydrogen and oxygen are the most im- 
portant. 

Whenever a compound containing carbon, hydrogen, and 
nitrogen is heated in a closed vessel, so that the air does 
not have access to it and it cannot burn up, the nitrogen 
passes out of the compound, not as nitrogen, but in com- 
bination with hydrogen, in the form of the compound 
called ammonia. Nearly all animal substances contain 
carbon, hydrogen, oxygen, and nitrogen, and many of them 
give off ammonia w r hen heated. Similarly, compounds con- 
taining carbon, oxygen, and hydrogen, even though they 
are thoroughly dry, when heated give off oxygen in com- 
bination w T ith hydrogen in the form of water. Some animal 
substances give off ammonia when they undergo decompo- 
sition in the air. The coal which is used for making 
illuminating gas contains some hydrogen and nitrogen in 
chemical combination, and when the coal is heated ammo- 
nia is given off. 

When animal substances undergo decomposition in the 
presence of a base w T here the temperature is comparatively 

133 



134 INTRODUCTION TO CHEMISTRY. 

high, the nitrogen combines with oxygen and the metal 
of the base. Either a nitrite or a nitrate is formed ; that 
is to say, either a salt of nitrous acid, HN0 2 , or of nitric 
acid, HN0 3 . In some countries where the conditions are 
favorable to the process, immense quantities of nitrates are 
found, chiefly potassium nitrate, or saltpetre, KN0 3 , and 
sodium nitrate, or Chili saltpetre, NaN0 3 . The change of 
the animal substances to the form of nitrates is probably 
caused by myriads of minute living organisms. How they 
effect the change is not known. From the salts of nitric 
acid which are found in nature, nitric acid itself can easily 
be prepared. 

Nearly all the compounds of nitrogen with which we 
shall have to deal are made either from ammonia or from 
nitric acid. 

Ammonia, NH 3 . — The conditions under which ammonia 
is formed have been mentioned. The chief source at pres- 
ent is the " ammonia-water " of the gas-works. This is 
the water through which the gas has been passed for the 
purpose of removing the ammonia, which passes into solu- 
tion. By adding hydrochloric acid to this liquid, am- 
monium chloride, which is a compound of the acid with 
ammonia, is formed. This is the well-known substance sal 
ammoniac. It appears that this name had its origin in the 
fact that common salt was formerly called sal armeniacum, 
and that afterward, through a misunderstanding, am- 
monium chloride came to be known by the same name 
which underwent change to the form sal ammoniacum, or 
sal ammoniac. 

As ammonium chloride, or sal ammoniac, is the most 
common compound containing ammonia, it is used in the 
laboratory for making ammonia. For this purpose it is 
only necessary to treat the salt with an alkali. 

Experiment 70. — To a little ammonium chloride on a watch- 
glass add a few drops of a strong solution of caustic soda, and 



PREPARATION OF AMMONIA. 135 

notice the odor of the gas given off. Do the same thing with 
caustic potash. Mix small quantities of ammonium chloride and 
lime in a mortar and notice the odor. Has the ammonium 
chloride itself this odor ? 

~pi eparation of Ammonia. — Ammonia is best prepared by 
mixing slaked lime and ammonium chloride. 

In addition to the ammonia, which is given off in the 
form of gas, calcium chloride, CaCl 2 , and water are formed 
in this reaction. It is represented thus: 

2NH 4 01 + Ca(OH), = 2NH 3 + CaCl 2 + 2H 2 0. 

This curious reaction will be more fully discussed after 
ammonia has been studied. 

Experiment 71. — Arrange an apparatus as shown in Figure 30, 
p. 106, omitting, however, the funnel-tube ; a cork with one open- 
ing will therefore serve. Weigh 100 grams quick-lime into a dish, 
add just enough water to slake it without making it moist, mix 
with 50 grams ammonium chloride, and transfer the whole to the 
flask. Push the stopper into place and gently heat the flask, 
which rests upon a sand-bath. After the air is driven out, the 
gas will be completely absorbed by the water in the first Wolff's 
flask. Disconnect the delivery-tube from the Wolff's flasks, and 
connect with another tube bent upward. Collect some of the 
escaping gas by displacement of air, placing the vessel with the 
mouth downward, as the gas is much lighter 
than air. The arrangement is shown in Fig. 
36. The vessel in which the gas is collected 
* should be dry, as water absorbs ammonia very 
readily. Hence, also, it cannot be collected over 
water. In the gas collected introduce a burning 
stick or taper. Ammonia does not burn in air, 
nor does it support combustion. In working 
with the gas great care must be taken to avoid 
inhaling it in any quantity. After enough has 
been collected in cylinders to exhibit the chief Fig. 36. 

properties, connect the delivery-tube again with the Wolff's 
flasks, and pass the gas over the water as long as it is given off. 




136 



INTRODUCTION TO CHEMISTRY. 



Properties of Ammonia. — From the observations made 
in the experiments just performed, it is seen that ammonia 
is a colorless, transparent gas. It has a very penetrating 
characteristic odor. In concentrated form it causes suffo- 
cation. Its specific gravity is 0.586; that is to say, it is but 
little more than half as heavy as air. It can easily be 
reduced to the liquid form by pressure and cold. When 
the pressure is removed from the liquefied ammonia, it 
passes back to the form of gas. In so doing it absorbs 
heat. These facts are taken advantage of in the artificial 
preparation of ice. Carre's ice-machine depends upon this 
principle. 

Ammonia does not burn in the air, but does burn in 
oxygen. It is absorbed by water in very large quantity. 
One volume of water at the ordinary temperature dissolves 
about 600 volumes of ammonia-gas, and at 0° about 1000 
volumes. 



[Problem. — A litre of air at 0° weighing 1.293 grams, and the 
specific gravity of ammonia gas being 0.586, how much would a 
litre of water increase in weight by being saturated with am- 
monia at 0° ?] 

The solution of ammonia in water is what is generally 
called ammonia in the laboratory. It is called " Spirits of 
Hartshorn" in common language. The solution has the 
odor of the gas. It loses all its gas when heated to the 
boiling-temperature. The solution shows a strong alkaline 
reaction and has the power to neutralize acids. 

Experiment 72. — Put 100 cc. dilute ammonia solution in an 
evaporating-dish. Try its effect on red litmus paper. Slowly 
add dilute hydrochloric acid until the alkaline reaction is de- 
stroyed and the solution is neutral. Evaporate to dryness on a 
water-bath. Compare the substance thus obtained with sal 
ammoniac or ammonium chloride. Taste them. Heat them on 
a piece of platinum-foil. Treat them with a caustic alkali. 
Treat with a little concentrated sulphuric acid in dry test-tubes. 






SALTS FORMED BY AMMONIA. 137 

Do they appear to be identical ?— The product is ammonium 
chloride, NH 4 C1. Similarly sulphuric acid and ammonia yield 
ammonium sulphate ; nitric acid and ammonia yield ammonium 
nitrate, etc. 

Experiment 73. — Fill a cylinder with ammonia-gas, and an- 
other of the same size with hydrochloric-acid gas. Bring them 
together with their mouths covered. Quickly remove the covers, 
when a dense white cloud will appear in and about the cylinders. 
This will soon settle on the walls of the vessels as a light white 
solid. It is ammonium chloride. Thus, from two colorless gases 
a solid substance is obtained by an act of chemical combination. 
Heat is evolved in the act of combination. 

Salts Formed by Ammonia. — It has been shown that the 
alkalies are strong bases, and that bases are compounds of 
metals with hydrogen and oxygen. Certainly those sub- 
stances which show an alkaline reaction are compounds of 
metals with hydrogen and oxygen. But in the solution of 
ammonia in water we have a substance which shows an 
alkaline reaction and which acts in nearly all respects very 
much like a solution of sodium hydroxide or potassium 
hydroxide. The salts which ammonia forms with acids 
are very similar to sodium and potassium salts. What is 
the substance which has the alkaline reaction ? and what 
are the salts which are formed by the action of acids on 
ammonia ? In the first place, it has been found that when 
an acid acts on ammonia the two combine directly without 
the formation of anything but the salt. Thus ammonia 
and hydrochloric acid form ammonium chloride : 

NH^ + HC1 = NH.C1. 

Ammonia and nitric acid form ammonium nitrate: 

NH 3 + HN0 3 = NH,N0 3 , etc., etc. 

Ammonium Theory. — On comparing the formulas of 
ammonium salts with those of potassium and sodium salts 
we see that, while in the potassium and sodium salts the 



138 INTRODUCTION TO CHEMISTRY. 

metals potassium and sodium take the place of the hydro- 
gen of the acids, in the ammonium salts the place of the 
hydrogen of the acid is taken by a compound of the for- 
mula NH 4 . It has been suggested, and the idea has gen- 
erally been accepted, that when ammonia-gas dissolves in 
water an unstable compound of the formula NH 4 OH is 
formed thus : 

NH 3 +H 2 = NH,OH. 

In this hydroxide, as in the salts of ammonia, the com- 
pound NH 4 appears to play the part of a metal. The 
compound NH 4 is, however, wholly hypothetical. As it 
appears to be this which plays the part of a metal in the 
solution as well as in the salts, tlie name ammonium has 
been given to it, the ending ium being that which is usu- 
ally given to signify metallic character. We speak, then, 
of ammonium salts, just as we speak of potassium or 
sodium or calcium salts. In the ammonium salts the 
hypothetical compound metal ammonium, NH 4 , is as- 
sumed to be present. If, however, we attempt to set it 
free or to set its hydroxide free, we get ammonia. On 
treating ammonium chloride with lime, if any action takes 
place at all, we should expect it to be that represented by 
the equation 

2NH 4 C1 + Ca(OH) 2 = CaCl 2 + 2NH.OH ; 

that is to say, we should expect the calcium and ammo- 
nium to exchange places,, Perhaps this is the action which 
takes place at first. But the compound NH 4 OH, or am- 
monium hydroxide, if formed at all, breaks up at once into 
ammonia and water, thus : 

NH.OH = NH 3 + H 2 0. 

So, too, if ammonium hydroxide, NH 4 OH, is present in 
the solution of ammonia in water, it breaks up very read- 



COMPOSITION OF AMMONIA. 139 

ily into ammonia and water under the influence of gentle 
heat, and ammonia gas is given off. 

Composition of Ammonia by Weight. — By oxidation 
under the proper conditions it is possible to convert the 
hydrogen of ammonia into water and leave the nitrogen in 
the free state. As water and nitrogen are the only prod- 
ucts formed, and the quantity of oxygen used up in the 
oxidation is equal to the quantity of oxygen found in the 
water formed, it follows that nitrogen and hydrogen are 
the only elements contained in ammonia. 

When electric sparks are passed for some time through 
a mixture of nitrogen and hydrogen, some ammonia is 
formed. Conversely, when electric sparks are passed for 
a time through ammonia, nitrogen and hydrogen are 
obtained. 

If, in the oxidation of a known quantity of ammonia, 
the water formed and the nitrogen left uncombined are ac- 
curately determined, it will be found that in ammonia the 
elements are combined in the proportion oi fourteen parts 
by weight of nitrogen to three parts by weight of hydrogen. 
This fact is expressed by the formula NH 3 ; 14 being the 
atomic weight of nitrogen. 

Composition of Ammonia by Volume. — The proportion 
by volume in which the two elements combine may be 
determined by the following method. When ammonia is 
treated with chlorine it is decomposed, the chlorine com- 
bining with hydrogen, and the nitrogen being left uncom- 
bined. The reaction is represented thus : 

NH 3 + 3C1 = N + 3HC1. 

Hydrogen and chlorine unite in equal volumes, as has 
already been seen. Now, if a solution of ammonia is 
added to a measured volume of chlorine until the chlorine 
is all used up, the volume of hydrogen which is extracted 



140 



INTRODUCTION TO CHEMISTRY. 



from ammonia is equal to the volume of chlorine used. 
The nitrogen left over was combined with the hydrogen 
whose volume has already been determined. It would be 
found that the volume of nitrogen is to that of the hydro- 
gen with which it was combined as 1 to 3 ; or in am- 
monia 1 volume of nitrogen is combined tvith 3 volumes of 
hydrogen. 

When a given volume of ammonia is decomposed into 
nitrogen and hydrogen, the mixture occupies just twice 
the volume that the ammonia did ; or, if a mixture of 
nitrogen and hydrogen in the proper proportions to form 
ammonia is caused to combine, the ammonia formed 
occupies one half the volume occupied by the mixture 
of gases. 

Relations between the Volumes of Combining Gases. — In 

studying the volume-relations of hydrogen, chlorine, and 
hydrochloric acid with reference to one another, we found 
that when hydrogen and chlorine combine one volume of 
the one combines with one volume of the other, and two 
volumes of the product are formed. These facts may be 
represented graphically thus : 



1 vol. 
hydrogen 



and combine to form 



1 vol. 
chlorine 



2 volumes hy- 
drochloric acid. 



When hydrogen and oxygen combine, two volumes of 
hydrogen combine with one volume of oxygen ; and the 
three volumes of gas thus combined form two volumes of 
water-vapor: 



VOLUME-RELATIONS BETWEEN COMBINING GASES. 141 
2 volumes hydrogen 



and combine to form 



1 vol. 
oxygen 



2 volumes 
water-vapor. 



Finally, it has just been shown that one volume of 
nitrogen combines with three volumes of hydrogen to 
form two volumes of ammonia: 

3 volumes hydrogen 



and 



combine to form 



1 vol. 
nitrogen 



2 volumes 
ammonia. 



A careful study of the volumes of combining gases has 
shown that these volumes always bear a simple relation to 
one another and to the volumes of the products formed. 
The three cases above presented show the more common 
relations met with among the elements. 

It is clear that the three elements chlorine, oxygen, and 
nitrogen influence hydrogen differently. One volume of 
chlorine can hold in combination but one volume of hy- 



142 INTRODUCTION TO CHEMISTRY. 

drogen. One volume of oxygen can hold in combination 
two volumes of hydrogen, and at the same time cause a 
condensation of volume from three volumes of gas to two. 
One volume of nitrogen can hold in combination three 
volumes of hydrogen, and, at the same time, cause the 
condensation of four volumes of gas to two. 

Relations between the Specific Gravities of Gases and 
their Atomic Weights. — Attention has already been called 
to the fact that the weights of equal volumes of hydrogen, 
chlorine, and oxygen stand in the same relation to one 
another as the combining weights. Nitrogen is no excep- 
tion to this rule. The specific gravity of nitrogen is 0.967. 
One litre of nitrogen weighs 1.2553 grams. The specific 
gravity of hydrogen is 0.0693; and the weight of a litre of 
hydrogen is 0.089578. But 0.0693 : 0.967 : : 1 : 14 and, of 
course, 0.089578 : 1.2553 : : 1 : 14. The accepted combin- 
ing weight of nitrogen is 14. 

These remarkable facts may be represented graphically 
thus : 



1 litre of hydro 0:1 1 litre of chlorine 1 litre of oxygen 1 litre of nitrogen 
weighs 0.089578 gr. weighs 3.17 gr. weighs 1.429 gr. weighs 1.2553 gr. 

These figures bear to one another the relations expressed 
by the figures 1, 35.4, 16, and 14. But these last figures 
very nearly express the atomic weights of the elements. 
It appears, therefore, that the atomic iveights of some of the 
gaseous elements bear to one another the same relations as 
the weights of equal volumes of the gases. 

Observations of this kind, together with other observa- 
tions on the conduct of gases, have led to a very important 
conception in regard to the nature of gases and the consti- 
tution of matter. This will be treated of farther on. For 



NITRIC ACID. 143 

the present it will be best to keep to the facts, so that 
before taking up speculations in regard to the hidden 
causes of the phenomena observed we may gain a solid 
foundation for these speculations. 

Nitric Acid, HN0 3 . — To effect the direct union of nitro- 
gen with oxygen and hydrogen is not easier than to effect 
the direct union of nitrogen with hydrogen to form ammo- 
nia. Nevertheless, the silent and continuous action of 
minute organisms in the soil is always tending to trans- 
form the waste-products of animal life into compounds 
closely allied to nitric acid. The process of nitrification 
has already been referred to. It is plainly an oxidizing 
process. In general, by oxidation the nitrogen of animal 
substances is converted into nitric acid, while by reduction 
it is converted into ammonia. 

Preparation of Nitric Acid. — In preparing nitric acid a 
nitrate is always used as the starting-point, and in this the 
metal is replaced by hydrogen. This is done in the same 
way that the metal sodium in sodium chloride is replaced 
by hydrogen in the preparation of hydrochloric acid, — viz., 
by treating the salt with a strong acid : 

2XaCl + H 2 S0 4 = Na.SO, + 2HC1; 
2XaX0 3 + H 2 S0 4 = Xa 2 S0 4 + 2HX0 3 . 

In these cases the acids obtained are volatile under the 
conditions of the experiment. Acids that are not vol- 
atile decompose the salts of those that are. At a suffi- 
ciently high temperature sulphuric acid is volatile, and an 
acid that is not volatile at such a temperature will set sul- 
phuric acid free. 

Experiment 74.— Arrange an apparatus as shown in Fig. 37. 
In the retort put 40 grams sodium nitrate (Chili saltpetre) and 
20 grams concentrated sulphuric acid. On gently heating, nitric 
acid will distil over, and be condensed in the receiver. After 



144 



INTRODUCTION TO CHEMISTRY, 



the acid is all distilled off, remove the contents of the retort. Re- 
crystallize the substance from water, and compare it with the 
sodium sulphate obtained in the preparation of hydrochloric acid. 
(See Experiment 58.) In the latter stage of the operation the 




Fig. 37. 



vessels become filled with a reddish-brown gas. 
is collected has a somewhat yellowish color. 



The acid which 



Pure nitric acid is a colorless liquid. It gives off color- 
less fumes when exposed to the air. When boiled it 
undergoes slight decomposition Into oxygen, water, and 
compounds of nitrogen and oxygen. One of these com- 
pounds is colored, and it is this that is noticed in the above 
experiment, and whenever strong nitric acid is boiled. 
Nitric acid undergoes a similar decomposition when exposed 
to the action of the direct rays of the sun. In consequence 
of this decomposition bottles containing strong nitric acid 
always contain a reddish-brown gas above the liquid after 
standing for some time. It acts violently on a great many 
substances, disintegrating them. It causes bad wounds in 
contact with the flesh j eats through clothing; burns wood; 



NITRIC ACID AN OXIDIZING AGENT. 145 

dissolves metals; and is altogether one of the most active 
of chemical substances. In working with the concentrated 
acid it is necessary to exercise the greatest care. 

Commercial nitric acid contains only about 68 per cent 
of the chemical compound HX0 3 . The rest is mostly 
water, though there are several impurities in small quan- 
tity. In order to get concentrated pure acid from this it 
must be distilled after the addition of some concentrated 
sulphuric acid. 

Experiment 75. — Mix together 200 grams concentrated sul- 
phuric acid and 100 grams ordinary concentrated nitric acid. 
Pour the sulphuric acid into the nitric acid. Distil the mixture 
slowly from a retort arranged as in Experiment 74, taking care to 
keep the neck of the retort cool by placing filter-paper moistened 
with cold water on it. Use the acid thus obtained for the purpose 
of studying the properties of pure nitric acid. 

Nitric Acid an Oxidizing Agent.— In consequence of the 
ease with which nitric acid decomposes, giving up oxygen, 
it is an excellent oxidizing agent, and is much used in the 
laboratory in this capacity. The following experiments 
illustrate this action: 

Experiment 76. — Pour [concentrated nitric acid into a wide 
test-tube, so that it is about one fourth filled. Heat the end of 
a stick of charcoal of proper size, and, holding the other end with 
a forceps, introduce the heated end into the acid. It will con- 
tinue to burn with a bright light, even though it is placed below 
the surface of the liquid. The action is oxidation. The char- 
coal in this case finds the oxygen in the acid, and not in the air. 
Great care must be taken in performing this experiment. The 
charcoal should not come in contact with the sides of the test- 
tube. A large beaker-glass should be placed beneath the test- 
tube, so that, in case it should break, the acid would be caught 
and prevented from doing harm. The arrangement of the ap- 
paratus is shown in Fig. 38. 

The gases given off from the tube are offensive and poisonous. 



146 



INTRODUCTION TO CHEMISTRY. 



Hence this as well as all other experiments with nitric acid should 
be carried on under a hood in which there is a good draught. 




Fig. 38. 



Experiment 77. — Boil a little strong nitric acid in a test-tube 
in the upper part of which some horse-hair has been introduced 
in the form of a stopper. The horse-hair will take fire and burn, 
leaving a white residue. Hold the test-tube with a forceps over 
a vessel to catch the contents should the tube break. 

Experiment 78. — In a small flask put a few pieces of granu- 
lated tin. Pour on this just enough ordinary concentrated nitric 
acid to cover it. Heat gently over a small flame. Soon action 
will take place. Colored gases will be evolved, the tin will dis- 
appear, and in its place will be found a white powder. This 
consists mostly of tin and oxygen. (See Experiment 15.) 

Action of Nitric Acid on Metals. — Like other acids, 
nitric acid forms salts with the metals. These can be 
made by treating the metals themselves with the acid, but 
in this case the formation of the salt is accompanied by 
another kind of action which is quite characteristic of 
nitric acid. The acid gives up a part of its oxygen and is 
thus converted into compounds of oxygen and nitrogen 



ACTION OF NITRIC ACID ON METALS. 147 

which contain a smaller proportion of oxygen than the 
acid. The compound most commonly formed in this way 
is nitric oxide, NO. It appears probable that the metal 
first abstracts oxygen from the acid, and that the oxide 
thus formed then dissolves in a part of the acid as the 
nitrate. The action in the case of copper should, accord- 
ing to this, be represented as taking place in two stages. 
First, nitric oxide, water, and copper oxide are formed as 
represented in this equation : 

(1) 2HN0, + 3Cu = H,0 + 3CuO + 2N0. 

Then the copper oxide forms copper nitrate with some of 
the acid, and this nitrate dissolves : 

(2) CuO + 2HN0 3 == Cu(N0 3 ) 2 + H 2 0. 

It is possible that to some extent hydrogen is set free 
from the acid by the action of the metal, and that this 
acts upon the acid, reducing it to lower oxides of nitrogen. 
Nitric oxide, NO, unites with oxygen from the air, and 
forms nitrogen peroxide, N0 2 , a colored gas, which is always 
seen when nitric acid acts upon metals. 

Experiment 79. —Dissolve a few pieces of copper foil in ordi- 
nary commercial nitric acid diluted with about half its volume of 
water. The operation should be carried on in a good-sized flask 
and under an efficient hood. When the copper has disappeared, 
pour the blue solution into an evaporating-dish, and evaporate 
down to crystallization. Compare the substance thus obtained 
with copper nitrate. — Heat a specimen of each. — Treat small 
specimens with sulphuric acid. — Do the two substances appear to 
be identical ? 

Experiment 80. — Heat specimens of potassium nitrate, sodium 
nitrate, lead nitrate, and any other nitrates which may be avail- 
able. All are decomposed giving off oxygen, in some cases mixed 
with oxides of nitrogen, among which is nitrogen peroxide, which 
can be recognized by its color. 



148 INTRODUCTION TO CHEMISTRY. 

General Properties of the Salts of Nitric Acid. — All salts 
of nitric acid are decomposed by heat, and all are soluble 
in water. 

Experiment 81. — Try the solubility in water of the nitrates 
used in the last experiment. 

Ammonia Formed by Reduction of Nitric Acid. — The 
formation of ammonia by reduction of nitric acid may be 
shown by the following experiment. 

Experiment 82. — In a good-sized test-tube treat a few pieces 
of granulated zinc with dilute sulphuric acid. What is evolved ? 
Now add drop by drop dilute nitric acid. Pour the contents of 
the tube into an evaporating-dish and evaporate the liquid. Put 
the residue into a test-tube and add caustic-soda solution. What 
is given off? Try the action of the gas on red litmus paper. 
Moisten the end of a glass rod with a little hydrochloric acid and 
hold it in the tube. What do you observe ? Explain it. Do the 
same with nitric acid. What are the fumes in this case ? 

Aqua Regia. — Aqua regia is made by mixing together 
concentrated nitric and hydrochloric acids. It received its 
name for the reason that it can dissolve gold, the king of 
metals. It is an excellent solvent, and is much used in 
the laboratory. 

Nitrous Acid, HN0 2 . — Among the reduction-products of 
nitric acid is nitrous acid, HN0 2 . This acid is most 
easily prepared in the form of a salt by reducing a nitrate. 
Thus, if potassium nitrate, KN0 3 , is melted together with 
metallic lead, the lead extracts a part of the oxygen and 
leaves potassium nitrite, KN0 2 . 

KN0 3 + Pb = KN0 2 + PbO. 

Experiment 83. — Heat together in a shallow iron plate 25 grams 
potassium nitrate and about 50 grams metallic lead. When both 
are melted stir them together as thoroughly as possible. After 



DECOMPOSITION OF NITRIC ACID. 149 

the mass is cooled down, break it up and treat with water in a 
flask. The potassium nitrite will dissolve, while the lead oxide 
and unused lead will not dissolve. Filter. Add a little sul- 
phuric acid to some of the solution. What is given off? See 
whether a solution of potassium nitrate acts in the same way. 

Nitrous Acid breaks down into Nitrogen Trioxide and 
Water. — When an acid is added to a solution of a nitrite, 
the salt is decomposed and nitrogen trioxide or nitrous 
anhydride, N 2 3 , is given off. Were the action in this 
case analogous to that which takes place when sulphuric 
acid acts upon a nitrate or upon a chloride, it w r ould be 
represented thus : 

2KN0 2 + H 2 S0 4 = K 2 S0 4 + 2HN0 2 . 

Nitrous acid would be formed; but instead of this a sub- 
stance which is nitrous acid less the elements of w r ater is 
formed : 

2HN0 2 = N 2 3 + H 2 0. 

This tendency on the part of compounds containing 
hydrogen and oxygen to decompose with formation of 
water is very commonly observed. We have already had to 
deal with a case of the kind in ammonium hydroxide. 
This substance, which probably exists in solution in water, 
yields ammonia and water when heated. Many compounds 
that do not break up in this way at ordinary temperatures 
do so at elevated temperatures. This decomposition is to be 
ascribed to the strong tendency of hydrogen to combine with 
oxygen. In complex compounds several forces are at work 
to keep the constituents in equilibrium. If the attraction 
of hydrogen for oxygen is much stronger than the other 
forces at work, the equilibrium is disturbed, and decom- 
position takes place. At least, this is the thought that 
naturally suggests itself by way of partial explanation of 
the phenomenon. 



150 



INTRODUCTION TO CHEMISTRY. 



Anhydrides. — A compound which, in its composition, 
bears to an acid the relation that nitrogen trioxide, N 2 3 , 
bears to nitrous acid, HN0 2 , is called an anhydride. 
Thus we have nitrous anhydride, N 2 3 ; nitric anhydride, 
N 2 B , etc. Nitric anhydride bears the same relation to 
nitric acid that nitrous anhydride bears to nitrous acid. 

N 2 3 + H 2 = 2HN0 2 ; 
N 2 6 + H 2 = 2HN0 3 . 

In more general terms, it may be said that any oxide 
which, when brought together with water, forms an acid 
by direct combination, is an anhydride. Other examples 
of this class of compounds will be met with farther on. 

The Oxides of Nitrogen, — Nitrogen and oxygen form 
five compounds with each other, of which all but one 
have already been mentioned. The names and for- 
mulas of the five compounds are nitrogen peroxide, N0 2 ; 
nitric oxide, NO; nitrous oxide, N 2 0; nitrous anhydride, 
N 2 3 ; and nitric anhydride, N 2 5 . If the formulas of 
these compounds are arranged in a series, beginning with 
that one which contains the smallest proportion of oxygen, 
it will be seen that the series affords a striking illustration 
of the facts from which the law of multiple proportions is 
deduced. The series is: 

Nitrous oxide N 2 

Nitric oxide NO or N a 2 

Nitrogen trioxide N 2 3 

Nitrogen peroxide N0 2 or N 2 4 

Nitric anhydride N 2 6 

It will be seen that the quantities of oxygen combined 
with 28 parts by weight of nitrogen are 16, 32, 48, 64, and 80. 

[What other series of compounds have you already had 
to deal with which illustrates the law of multiple propor- 
tions almost equally strikingly ?] 



NITROUS OXIDE. 151 

Of the oxides of nitrogen, only three need be studied 
here, and after what has already been said they need be 
studied only briefly. 

Nitrous Oxide, N 2 0. — This compound is formed by reduc- 
tion of nitric acid when the acid acts upon metals and the 
degree of concentration and the temperature are favorable. 
It is usually prepared by heating ammonium nitrate, 
NH 4 N0 3 . The decomposition takes place as represented 
thus : 

NH 4 NO, = N 2 -f 2H 2 0, 

the products being nitrous oxide and water. In this reac- 
tion the tendency of hydrogen and oxygen to combine at 
elevated temperatures is shown. At ordinary temperatures 
this tendency is not strong enough to cause a disturbance 
of the equilibrium of the parts of the compound. As the 
temperature is elevated it becomes stronger and stronger, 
until finally the decomposition above represented takes 
place and the elements combine. 

Experiment 84.— In a retort heat 10 to 15 grams crystallized 
ammonium nitrate until it has the appearance of boiling. Do 
not heat higher than is necessary to secure a regular evolution of 
gas. Connect a wide rubber tube directly with the neck of the 
retort and collect the evolved gas over water, as in the case of 
oxygen. It supports combustion almost as well as pure oxygen. 
Try experiments with wood, a candle, and a piece of phosphorus 
in a deflagrating-spoon. 

Properties. — The gas is colorless and transparent. It 
has a slightly sweetish taste. It is somewhat soluble in 
water, so that when collected over water there is always 
considerable loss. When inhaled it causes a kind of in- 
toxication, which is apt to show itself in the form of hys- 
terical laughing, hence the name laughing-gas. Inhaled 
in larger quantity it causes unconsciousness and insensi- 



152 INlRODtJCTlON TO CHEMISTRY, 

bility to pain. It is therefore used to prevent pain in 
minor surgical operations, as, for example, in pulling 
teeth. 

Nitrous oxide is easily converted into a liquid by cold and 
pressure. In this form it can now be bought contained in 
strong iron cylinders. On opening the stop-cock of the 
cylinder the substance escapes in gaseous form. 

Nitric Oxide, NO. — This gas is formed when nitric acid 
acts upon some metals, as copper. The action is believed 
to involve two changes, as described on p. 147. 

The two equations representing the action may be com- 
bined in one thus : 

8HNO3 + 3Cu = 3Cu(N0 3 ) 2 + 4H 2 + 2NO. 

As already stated, however, it is possible that the for- 
mation of the nitric oxide is due to the two reactions 
represented in the equations : 

2HN0 3 + Cu = Cu(N0 3 ) 2 + 2H, and 
2HN0 3 + 6H = 4H 2 + 2NO. 

Experiment 85.— Arrange an apparatus as shown in Fig. 39. 
In the flask put a few pieces of copper foil. Cover 
this with water. Now add slowly, waiting each 
time for the action to begin, ordinary concentrated 
nitric acid. When enough nitric acid has been 
added gas will be evolved. If too much acid is 
added, it not unfrequently happens that the evolu- 
tion of gas takes place too rapidly, so that the 
liquid is forced out of the flask through the funnel- 
tube. This can be avoided by not being in a hurry. 
At first the vessel becomes filled with a reddish- 
j| brown gas, but soon the gas evolved becomes color- 
less. Collect over water two or three vessels full. 
The gas collected is principally nitric oxide, NO, 
Fig. 39. though it is frequently mixed with a considerable 
quantity of nitrous oxide. 

Experiment 86.— Turn one of the vessels containing colorless 




NITRIC OXIDE. 153 

nitric oxide with the mouth upward and uncover it. A colored 
gas is at once seen, presenting a very striking appearance. Do 
not inhale the gas. Perform the experiments with nitric oxide 
where there is a good draught. 

Properties of Nitric Oxide. — Nitric oxide is a colorless, 
transparent gas. Its most remarkable property is its 
power to combine directly with oxygen when the two are 
brought together. The act of combination is not accom- 
panied by the appearance of light, though heat is evolved. 
The reaction is represented by the equation 

NO + = N0 2 . 

The product is nitrogen peroxide, and this at ordinary 
temperatures is a reddish-brown gas. 

Nitric oxide does not burn and does not support com- 
bustion. When the fact is borne in mind that nitrous 
oxide, N 2 0, supports combustion almost as well as oxygen, 
it appears strange that another compound of nitrogen and 
oxygen, containing twice as much oxygen relatively to the 
same quantity of nitrogen, should not support combustion. 
This is explained by the relative stability of the two com- 
pounds. In the case of nitrous oxide, the oxygen is not 
firmly held in combination; the equilibrium established 
between the forces at work is not a stable one. Hence, 
when a substance which readily combines with oxygen is 
brought in contact with it, the equilibrium is disturbed, 
or the oxide is decomposed. On the other hand, in nitric 
oxide the arrangement of the parts is a more stable one. 
The oxygen, although present in larger quantity than in 
nitrous oxide, is held more firmly, and cannot easily be 
extracted. The gas does not support combustion. 

Nitrogen Peroxide, N0 2 . — This gas is made by direct 
combination of nitric oxide with oxygen, as seen in the 
last experiment. It has a disagreeable smell and is poison- 



154 INTRODUCTION TO CHEMISTRY. 

ous. It gives up a part of its oxygen quite easily, and is 
hence useful as an oxidizing agent. 

Use of the Oxides of Nitrogen in the Manufacture of Sul- 
phuric Acid. — The higher oxides of nitrogen, especially the 
trioxide, N 2 3 , and the peroxide, N0 2 , readily give up 
oxygen, and are changed to nitric oxide, NO. If air is 
present, nitric oxide is changed back again to the higher 
oxides, which may again give up oxygen, again yielding 
nitric oxide, and so on indefinitely. It will thus be seen 
that these oxides of nitrogen may be made to serve the 
purpose of transferring oxygen from the air to other sub- 
stances. Advantage is taken of these facts in the manu- 
facture of sulphuric acid. 

Summary. — The simpler nitrogen compounds are made 
either from ammonia or from nitric acid. Ammonia is 
formed in nature by the spontaneous decomposition of 
animal substances. It is also formed by heating substances 
which contain carbon, hydrogen, and nitrogen. The prin- 
cipal source is the "ammonia-water" of the gas-works. 

Nitric acid is formed in nature as the potassium or 
sodium salt, by the action of certain organisms on sub- 
stances containing nitrogen. 

Ammonia is prepared from an ammonium salt by treat- 
ing it with a strong base. Ammonium chloride and lime 
are commonly used. 

With acids ammonia forms salts which are known as 
ammonium salts, and in which the compound NH 4 is sup- 
posed to act the part of a metal. This hypothetical metal 
is called ammonium. 

Ammonia consists of 14 parts of nitrogen to 3 parts of 
hydrogen. The gases are combined in the proportion of 1 
volume of nitrogen to 3 volumes of hydrogen. The 4 
volumes thus combined condense to 2 volumes of am- 
monia. 



SUMMAMT. 155 

There is always a simple relation between the volumes 
of combining gases and the volume of the compound 
formed if it is a gas. 

A comparison 6i the specific gravities of the gaseous 
elements shows that these bear to one another the same 
relation as the atomic weights. 

Nitric acid is prepared from a nitrate by treating it with 
sulphuric acid. It is comparatively unstable, giving up 
oxygen easily. With metals it yields salts, but the action 
involves the reduction of a part of the acid, and leads to 
the formation of various products, among which may be 
nitrous oxide, N 2 0; nitric oxide, NO; nitrous anhydride, 
N 2 3 ; and nitrogen peroxide, N0 2 . Under some circum- 
stances, the action may even go far enough to form 
ammonia. Nitrous acid itself is unstable, breaking up into 
the anhydride, N 2 3 , and water. 

Anhydrides are substances which, when brought together 
with water, combine with it to form acids. 

Though nitrous oxide is formed by reduction of nitric 
acid, it is best prepared in pure condition by heating 
ammonium nitrate. It supports combustion well. 

Nitric oxide is made by reduction of nitric acid by 
means of copper. It combines directly with oxygen, form- 
ing the strongly-colored and disagreeable-smelling nitrogen 
peroxide. 

Nitrogen peroxide gives up a part of its oxygen easily, 
and is hence a good oxidizing agent. It is thus reduced 
to nitric oxide, which in the air takes up oxygen. 



CHAPTER XI. 

CARBON. 

Carbfln in Plants and Animals. — Most substances of 
vegetable or animal origin blacken when they are heated 
to a sufficiently high temperature, and if heated in the air 
they burn up, as we say. This is due to the fact that 
nearly all animal and vegetable substances contain the ele- 
ment carbon. When they are heated the other elements 
present are first driven off in various forms of combination, 
while the carbon is the last to go. If the heating is carried 
on in the air, the carbon finally combines with oxygen to 
form a colorless gas — it burns up. Carbon is the central 
element of organic nature. There is not a living thing, 
from the minutest microscopic animal to the mammoth, 
from the moss to the giant tree, which does not contain 
this element as an essential constituent. The number of 
the compounds which it forms is almost infinite, and they 
present such peculiarities that they are commonly treated 
under a separate head, " Organic Chemistry." There is 
no good reason for this, except the large number of the 
compounds. For the present it will suffice to study the 
chemistry of the element itself, and of a few of its simpler 
compounds, and farther on a few chapters on some of its 
most important compounds will be presented. 

Occurrence. — From what has already been said, it will be 
seen that carbon occurs in nature for the most part in com- 
bination with other elements. It occurs not only in living 

156 



DIAMOND. 157 

things, but in their fossil remains, as in coal. Coal-oil, or 
petroleum, consists of a large number of compounds which 
contain only carbon and hydrogen. Most products of 
plant-life contain the elements carbon, hydrogen, and oxy- 
gen. Among the more common of these may be mentioned 
sugar, starch, cellulose, etc. Most products of animal life 
contain carbon, hydrogen, oxygen, and nitrogen. Among 
them may be mentioned albumin, fibrin, casein, the fats, 
etc. Carbon occurs in the atmosphere in the form of car- 
bon dioxide. [What evidence have you had of the presence 
of carbon dioxide in the air ?] It also occurs in the form 
of salts of carbonic acid — the carbonates, which are widely 
distributed, forming whole mountain-ranges. Limestone, 
marble, and chalk are calcium carbonate. 

Uncombined, the element occurs pure in two very differ- 
ent forms in nature: (1) as diamond; and (2) as graphite, 
or plumbago. 

Before presenting the evidence which leads to the con- 
clusion that diamond and graphite are only modifications 
of the same element, and that while closely related to each 
other they are also equally closely related to charcoal, it 
will be best to study separately the properties of each of 
these three substances. 

1. Diamond. — The diamond is found in but few places 
on the earth. Practically nothing is known as to the con- 
ditions which gave rise to its formation. The celebrated 
diamond-beds are in India, Borneo, Brazil, and South 
Africa. When found, diamonds are covered with an 
opaque layer, which must be removed before the beau- 
tiful properties appear. The crystals are sometimes what 
are known as octahedrons; that is to say, they are regular 
eight-sided figures, though usually they are somewhat more 
complicated. It is the hardest substance known. 

If heated to a very high temperature without access of 
air^ it swells up and is converted into a black mass re- 



158 INTRODUCTION TO CHEMISTRY. 

sembling graphite. This change takes place without loss 
in weight. Heated to a high temperature in oxygen, it 
burns up, yielding only carbon dioxide. It is insoluble in 
all ordinary liquids. 

2. Graphite.— Graphite, or plumbago, is found in nature 
in large quantities. Sometimes it is crystallized, but in 
forms entirely different from those assumed by the dia- 
mond. It can be prepared artificially by dissolving char- 
coal in molten iron, from which solution graphite is de- 
posited on cooling. It has a grayish-black color and a 
metallic lustre. It is quite soft, leaving a leaden-gray mark 
on paper when drawn across it, and is hence used in the 
manufacture of so-called lead pencils. It is sometimes 
called black-lead. 

When heated without access of air it remains unchanged. 
Heated to a very high temperature in the air, or in oxygen, 
it burns up, forming only carbon dioxide. Like the dia- 
mond, it is insoluble in all ordinary liquids. 

3. Amorphous Carbon. — All forms of carbon that are 
not diamond or graphite are included under the name 
amorphous carbon. The name signifies simply that it is 
not crystallized. The most common form of amorphous 
carbon is ordinary charcoal. 

Charcoal is that form of carbon which is made by tne 
charring process, which consists simply in heating without 
a free supply of air. The substance almost exclusively 
used in the manufacture of charcoal is wood. As has al- 
ready been stated, wood is made of a large number of sub- 
stances, nearly all of which, however, consist of the three 
elements carbon, hydrogen, and oxygen. One of the chief 
constituents of all kinds of wood is cellulose. Now, when 
a piece of wood is heated to the kindling temperature, it 
burns, if air is present. The chemical changes which take 



CHARCOAL. 159 

place are complex under ordinary circumstances; but if 
care is taken, the combustion can be made complete, when 
all the carbon is converted into carbon dioxide, and all the 
hydrogen into water. If, on the other hand, the air is pre- 
vented from coming in contact with the wood in sufficient 
quantity to effect complete combustion, the hydrogen is 
given off partly as water and partly in the form of volatile 
compounds containing carbon and oxygen. Most of the 
carbon, however, is left behind as charcoal, as there is not 
enough oxygen to convert it into carbon dioxide. 

A Charcoal-kiln. — A charcoal-kiln consists essentially of 
a pile of wood so arranged as to leave spaces between the 
pieces. The pile is covered with some rough material 
through which the air will not pass easily, as, for example, 
a mixture of powdered charcoal, turf, and earth. Small 
openings are left in this covering so that after it is kindled 
the wood will continue to burn slowly. The changes above 
mentioned take place, the gases or volatile substances pass- 
ing out at the top of the kiln, and appearing as a thick 
smoke. In due time the holes through which the air gains 
access to the wood, and which also make the burning 
possible, are closed, and the burning stops. Charcoal, 
which is impure amorphous carbon, is left behind. As 
wood always contains some incombustible substances in 
small quantity, these are, of course, found in the charcoal. 
When the wood or charcoal is burned, these substances re- 
main behind as the ash. 

Wood-charcoal. — Ordinary charcoal is a black, compara- 
tively soft substance. It burns in the air, though not 
easily unless the gases formed are constantly removed and 
fresh air is supplied, as when the draught is good or a pair 
of bellow r s is used. It burns readily in oxygen (see Experi- 
ment 22). The product of the combustion in oxygen and 
in air, when the conditions are favorable, is carbon dioxide, 



160 INTRODUCTION TO CHEMISTRY. 

C0 2 . In the air, when the draught is bad, another com- 
pound of carbon and oxygen, carbon monoxide, CO, is 
formed. Heated without access of air, charcoal remains 
unchanged. Charcoal is insoluble in all ordinary liquids. 

Coke. — Besides wood-charcoal, there are other forms of 
amorphous carbon, which are manufactured for special 
purposes, or are formed in processes carried on for the 
sake of other products. Colce is a form of amorphous car- 
bon which is made by heating ordinary gas-coal without 
access of air, as is done on a large scale in the manufacture 
of illuminating gas. Coke bears to coal about the same re- 
lation that charcoal bears to wood. 

Lamp-black is a very finely divided form of charcoal 
which is deposited on cold objects placed in the flames of 
burning oils. The oils consist almost exclusively of carbon 
and hydrogen. When burned in the air they yield carbon 
dioxide and water. If the flame is cooled down by any 
means, or if the supply of air is partly cut off, the carbon 
is not completely burned; the flame "smokes," and deposits 
soot : this soot is largely made up of fine particles of carbon. 
It is used in the manufacture of printer's ink. Carbon is 
acted upon directly by very few substances, and is not 
soluble, so that it is impossible to destroy the color of 
printer's ink without destroying the material upon which it 
is impressed. 

Bone-black, or Animal Charcoal, is a form of amorphous 
carbon which is made by charring bones. Unless treated 
with an acid it contains the incombustible substances con- 
tained in bone, as calcium phosphate, etc. 

Charcoal Filters. — Bone-black and wood-charcoal are 
very porous and have the power to absorb gases. When 
placed in air containing bad-smelling gases, these are 
absorbed and the air thus purified. When water contain- 



CHARCOAL FILTERS. 161 

ing disagreeable substances is treated with charcoal, these 
are wholly or partly absorbed and the water improved. 
Charcoal filters are therefore extensively used. A charcoal 
filter to be efficient should be of good size, and from time 
to time the charcoal should be renewed. The small filters 
which are screwed into faucets are of little value, as the 
charcoal soon becomes charged with the objectionable 
material of the water. 

Some coloring matters can be removed from liquids by 
passing the latter through bone-black filters. On the large 
scale, this fact is taken advantage of in the refining of 
sugar. The solution of sugar first obtained from the cane 
or beet is highly colored ; and if it were evaporated, the 
sugar deposited from it would be dark-colored. If, how- 
ever, the solution is first passed through bone-black filters, 
the color is removed, and now, on evaporating, white sugar 
is deposited. In the laboratory constant use is made of this 
method for the purpose of purifying liquids. 

Experiment 87. — Make a filter of bone-black by fitting a paper 
filter into a funnel 12 to 15 mm. (5 to 6 inches) in diameter at its 
mouth. Half-fill this with bone-black. Make a dilute solution 
of indigo. Pour it through the filter. If the conditions are right 
the solution will pass through colorless. — Do the same thing with 
a dilute solution of litmus. — If the color is not completely re- 
moved by one filtering, filter it again— The color can also be 
removed from solutions by putting some bone-black into them 
and boiling for. a time. — Try this with half a litre each of the 
litmus and indigo solutions used in the first part of the experi- 
ment. Use about 4 to 5 grams bone-black in each case. Shake 
the solution frequently while heating. 

Wood is Charred to Preserve it. — Charcoal does not 
undergo decay in the air or under water nearly as readily 
as wood. That is another way of stating the chemical fact 
that the substances of which wood is made up are more 
susceptible to the action of other chemical substances than 
charcoal is. The relative ease with which charcoal and 



162 INTRODUCTION TO CHEMISTRY. 

wood burn in the air illustrates this fact. Piles driven 
below the surface of water are charred to protect them from 
the action of those substances which cause decay. Most of 
the houses in Venice stand on piles of wood which have 
been charred. Oak stakes have quite recently been found 
in the Thames where, according to Tacitus, the Britons 
fixed stakes to prevent the passage of Caesar and his army. 

Coal. — Under this head are included a great many kinds 
of impure amorphous carbon which occur ready-formed in 
nature. Although an almost infinite number of kinds of 
coal are found, for ordinary purposes they are classified as 
hard and soft coals, or anthracite and bituminous coals. 
Then there are substances more nearly allied to wood 
called lignite, and those which represent a very early stage 
in the process of coal-formation, as, for example, peat. 

A close examination of all these varieties has shown that 
they have been formed by the gradual decomposition of 
vegetable material in an insufficient supply of air. The 
process has been going on for ages. Sometimes the sub- 
stances have, at the same time, been subjected to great 
pressure, as can be seen from the position in which they 
occur in the earth. The products in the earlier stages of 
the coal-forming process are, naturally, more closely allied 
to wood than those in the later stages. 

All forms of coal contain other substances in addition to 
the carbon. The soft coals are particularly rich in other 
substances. When heated they give off a mixture of gases 
and the vapors of volatile liquids. The gases are, for the 
most part, useful for illuminating purposes. The liquids 
form a black, tarry mass known as coal-tar, from which are 
obtained many valuable compounds of carbon. The gases 
are passed through water for the purpose of removing cer- 
tain impurities. This water absorbs ammonia and forms 
the ammonia-water of the gas-works, which, as has been 
stated, is the principal source of ammonia. 



THREE FORMS OF CARBON. 163 

Diamond, Graphite, and Charcoal Different Forms of the 
Element Carbon. — An element is a kind of matter from 
which no simpler kind of matter can be obtained by any 
means now known to chemists. From hydrogen nothing 
but hydrogen can be obtained, except by bringing it 
together with some other element; from nitrogen, nothing 
but nitrogen, etc. In the case of carbon, however, the 
element appears in three forms which differ markedly from 
one another. It is difficult to conceive that the soft, black 
charcoal and the dull, gray, soft graphite are chemically 
identical with the hard, transparent, brilliant diamond. 
Yet this is undoubtedly the case, as can be % proved by a 
very simple experiment. Each of the substances when 
burned in oxygen yields carbon dioxide. Now, the com- 
position of carbon dioxide is known, so that if the weight 
of carbon dioxide formed in a given experiment is known, 
the weight of carbon contained in it is also known. When 
a gram of pure charcoal is burned it yields 3f grams of 
carbon dioxide, and in this quantity of carbon dioxide 
there is contained exactly 1 gram of carbon. Further, 
w T hen a gram of graphite is burned the same weight 
(3| grams) of carbon dioxide is formed as in the case of 
charcoal; and the same is true of diamond. It follows 
from these facts that the three forms of matter known as 
charcoal, graphite, and diamond consist only of the element 
carbon. 

[Problem. — How much carbon dioxide, C0 2 , should be ob- 
tained by burning 0. 5 gram diamond ? The atomic weight of 
carbon is 12.] 

Other Examples of the Occurrence of a Substance in 
Different Forms. — That one and the same substance can 
appear in markedly different forms under different con- 
ditions is seen in the case of water. Hail and snow would 
hardly be suspected of being the same substance by one 
who was not quite familiar with them. The difference in 



164 INTRODUCTION TO CHEMISTRY. 

this case, as in that under discussion, is believed to be due 
to the way in which the small particles of which the sub- 
stances are made up are arranged with reference to one an- 
other. If we had a number of small pieces of wood all of 
the same size and shape, say cubical, and should carefully 
arrange these in some regular way, we might easily make a 
comparatively compact mass of them, and the mass would 
have a regular form. We might arrange them, further, in 
a second way with regularity. And we might simply throw 
the pieces together in a jumble. These three kinds of 
arrangement would represent, in a rough way, the differ- 
ence between the three forms of carbon. Each pile would 
be made of wood, but still in outward appearance they 
would differ from one another. 

Common Properties of the Three Forms of Carbon. — Not- 
withstanding the marked differences in their appearance, 
the three forms of carbon have, as we have seen, some 
properties in common. They are insoluble in all ordinary 
liquids. They are tasteless and inodorous. They are 
infusible. When heated without access of air, they remain 
unchanged, unless the temperature is very high. 

Chemical Conduct of Carbon. — At ordinary temperatures 
carbon is an inactive element. If it is left in contact 
with any one of the elements thus far studied, — viz., hy- 
drogen, oxygen, chlorine, and nitrogen, — no change takes 
place. Indeed, unless the temperature is raised it will not 
combine with any other element. At higher tempera- 
tures, however, it combines with other elements, especially 
with oxygen, with great ease. Under proper conditions 
it combines also with nitrogen, with hydrogen, and with 
many other elements. It combines with oxygen either 
directly, as when it burns in the air or in oxygen ; or it 
abstracts oxygen from some of the oxides, 



DIRECT UNION OF CARBON AND OXYGEN 165 

Direct Union of Carbon and Oxygen. — This has already 
been illustrated in Experiment 22, and is also illustrated in 
every charcoal furnace. That carbon dioxide is the prod- 
uct may be shown by passing the gas into lime-water or 
baryta-water, when insoluble calcium, or barium, carbonate 
will be thrown down. 

Experiment 88. — Put a small piece of charcoal in a hard-glass 
tube. Pass oxygen through the tube, at the same time heatiug 
it. Pass the gases into clear lime-water. Arrange the appa- 
ratus as shown in Pig. 40. 




Fig. 40. 



A is a large bottle containing oxygen ; B is a cylinder contain- 
ing sulphuric acid ; C is a U-tube containing calcium chloride ; 
D is the hard-glass tube containing the charcoal ; E is the cylin- 
der with clear lime-water. In what previous experiment was 
this method of showing the formation of carbon dioxide used ? 
The reason why it is used is simply that an insoluble compound 
is formed, and this can be seen, and it can be separated from the 
liquid and examined. The reaction is represented thus : 



Ca(OH) 2 + C0 2 

Lime. Carbon dioxide. 



CaCOs + H 2 0. 
Calcium carbonate. 



166 



INTRODUCTION TO CHEMISTRY. 



No other common gas acts in this way on lime-water. Hence, 
when, under ordinary circumstances, a gas is passed into lime- 
water and an insoluble substance is formed, we may conclude 
that the gas is carbon dioxide. 

Abstraction of Oxygen from Compounds by Means of Car- 
bon. — This can be illustrated in a number of ways. 

Experiment 89. — Mix together 2 or 3 grams powdered copper 
oxide, CuO, and about one tenth its weight of powdered char- 
coal ; heat in a tube to which is fitted 
an outlet-tube, as shown in Fig. 41. 

Pass the gas which is given off into 
lime-water contained in a test-tube. 
Is it carbon dioxide ? What evidence 
have you that oxygen has been ex- 
tracted from the copper oxide ? What 
is the appearance of the substance 
left in the tube ? Does it suggest the 
metal copper ? Treat a little of it 
with strong nitric acid. What should 
take place if the substance is metallic 
copper? (See Experiment 79.) What 
does take place? The reaction between the charcoal and the 
copper oxide is represented thus : 

2CuO + C = 2Cu + C0 2 . 

Experiment 90. — Perform a similar experiment with a little 
white arsenic in a small glass tube closed at one end. Take 
about equal parts of charcoal and arsenic. White arsenic is a 
compound of the element arsenic and oxygen, of the composition 
represented by the formula As 2 3 . The reaction which takes 
place when it is heated with charcoal is represented thus : 

2As 2 3 + 3C = 4As + 3C0 2 . 




Fig. 41. 



The element arsenic is volatile, and is hence driven upward 
and deposited on the inside of the tube above the mixture, in the 
form of a mirror with metallic lustre. 



DEDUCTION. 167 

Reduction. — The abstraction of oxygen from a com- 
pound is known as reduction, as has already been ex- 
plained. Hence carbon is called a reducing agent. It is 
indeed the reducing agent that is most extensively used 
in the arts. Its chief use is in extracting metals from 
their ores. Thus, iron does not occur in nature as iron, 
but in combination with other elements, especially with 
oxygen. In order to get the metal, the ore must be 
reduced, or, in other words, the oxygen must be extracted. 
This is invariably accomplished by heating it with some 
form of carbon, either charcoal or coke. 

[What other element already studied acts as a reducing 
agent ? Give an example of its reducing power.] 



CHAPTEE XII. 

SOME OF THE SIMPLER COMPOUNDS OF CARBON. 

Compounds of Carbon with Hydrogen. — Carbon combines 
with hydrogen in a great many different proportions, the 
compounds being known as hydrocarbons. Among the 
simpler examples of these are marsh-gas, or fire-damp, 
CH 4 ; ethylene, 2 H 4 ; acetylene, C 2 H 2 ; and benzene, C 6 H 6 . 
These together with a few others will be taken up in a 
later chapter. 

Carbon Dioxide, C0 2 . — The principal compound of car- 
bon and oxygen is carbon dioxide, C0 2 , commonly called 
carbonic-acid gas. Under the head of The Atmosphere 
attention was called to the fact that this gas is a constant 
constituent of the air, though it is present in relatively 
very small quantity. It issues from the earth in many 
places, particularly in the neighborhood of volcanoes. 
Many mineral waters contain it in large quantity, as the 
waters of Pyrmont, Selters, and the Geyser Spring at Sara- 
toga. In small quantity it is present in all natural waters. 
In combination with bases it occurs in enormous quantities, 
particularly in the form of calcium carbonate, CaC0 3 , 
varieties of which are ordinary limestone, chalk, marble, 
and calc spar. Dolomite, a compound consisting of cal- 
cium carbonate and magnesium carbonate, MgC0 3 , enters 
largely into the structure of some mountain-ranges, as, for 
example, the Swiss Alps. 

168 



CARBON DIOXIDE. 169 

Carbon Dioxide given off from the Lungs. — Carbon dioxide 
is constantly formed in many natural processes. Thus, all 
animals that breathe in the air give off carbon dioxide from 
their lungs. 

Experiment 91.— Force the gases from the lungs through some 
lime-water by means of an apparatus arranged as shown in Fig. 
42. What is formed ? Add a few drops of hydrochloric acid. 
What takes place ? 

Carbon Dioxide formed in 
Combustion, in Decay, and in 
Fermentation. — That carbon 
dioxide is formed in the com- 
bustion of charcoal and wood 
has already been show T n. In a 
similar way it can be shown 
that the gas is formed when- 
ever any of our ordinary com- 
bustible materials are burned. Fm. 42. 
From our fires, as from our lungs, and from the lungs of 
all animals, then, carbon dioxide is constantly given off. 
Further, the natural processes of decay of both vegetable 
and animal matter tend to convert the carbon of this matter 
into carbon dioxide, which is then spread through the air. 
The process of alcoholic fermentation, and some other like 
processes, also give rise to the formation of carbon dioxide. 
In all fruit-juices there is contained sugar. When the 
fruits ripen, fall off, and undergo spontaneous change, the 
sugar is changed to alcohol and carbon dioxide. 

It is clear, therefore, that there are many important 
sources of supply of carbon dioxide, and it will readily be 
understood why the gas should be found everywhere in the 
air. 

Decomposition of Carbonates by Acids. — The easiest way 
to get carbon dioxide unmixed with other substances is to 




170 INTRODUCTION TO CHEMISTRY. 

add an acid to a carbonate. "Whenever an acid is added to 
a carbonate there is an evolution of gas. 

< Experiment 92. — In test-tubes add successively dilute hydro- 

^ chloric, sulghiirip, nitric, and acetic acids to a little sod ium re ir- 1 \ 

•^^onate^ Tneach case pass the £a*s given off through lime-wktefr^ 
and insert a burning stick in the upper part of each tube. — Per- 
form the same experiment with small pieces of marble. 

Comparison of this Decomposition with other similar 
Acts. — The decomposition of the salts of carbonic acid by 
other acids illustrates the principle that is involved in 
setting nitric acid free from a titrate, or hydrochloric acid 
from a chloride, by means of sulphuric acid. The non- 
volatile acid drives out the volatile acid. When, for exam- 
ple, hydrochloric acid is added to sodium carbonate the 
first action consists in an exchange of the hydrogen of the 
acid for the metal of the carbonate : 

Na 2 C0 3 + 2HC1 = 2NaCl + H 2 C0 3 . 

If sulphuric acid is used, the reaction is represented thus : 

Na.CO, + H 2 S0 4 = Na 2 S0 4 + H 2 C0 3 . 

This reaction is analogous to that which takes place be 
tween sodium nitrate and sulphuric acid in the preparation 
of nitric acid : 

2NaN0 3 + H 2 S0 4 = Na 2 S0 4 + 2HN0 3 . 

Carbonic acid, however, is an unstable substance, and 
breaks up into water and carbon dioxide as soon as it is 
liberated from its salts : 

H 2 C0 3 = H a O + C0 2 . 

As the carbon dioxide is a gas it is given off as a result of 
the decomposition. 

It will be seen that the decomposition of carbonic acid 



CARBON DIOXIDE, 



171 



into carbon dioxide and water is analogous to the decompo- 
sition of nitrous acid, HN0 2 , into nitrogen trioxide and 
water; and similar to the decomposition of ammonium 
hydroxide into ammonia and water. 

[What is a compound called which bears to an acid the 
relation that carbon dioxide bears to carbonic acid ?] 

Preparation of Carbon Dioxide.— For the purpose of pre- 
paring carbon dioxide in the laboratory, calcium carbonate in 
the form of marble, or limestone, and hydrochloric acid are 
commonly used. The reaction involved is represented thus : 

CaC0 3 + 2HC1 = CaCl 2 + C0 2 + H 2 0. 



Experiment 93. — Arrange an apparatus as shown in Fig. 43. 
In the flask put some pieces of marble, or lime- 
stone, and pour ordinary hydrochloric acid on 
it. The gas should be collected by displacement 
of air, the vessel being placed with the mouth 
upward, as the gas is much heavier than air. 
Collect several cylinders or bottles full of the 
gas. Into one introduce successively a lighted 
candle, a burning stick, a bit of burning phos- 
phorus in a deflagrating-spoon. What takes 
place ? With another proceed as if pouring water 
from it. Pour the invisible gas upon the flame 
of a burning candle. Pour some of the gas from 
one vessel to another, and show that it has been 
transferred. Weigh a beaker on a balance, and 
pour carbon dioxide into it. Give an account of ' 43 * 

the results obtained and state the conclusions you are justified 
in drawing from what you have seen. 




Physical Properties of Carbon Dioxide. — Carbon dioxide 
is a colorless gas at ordinary temperatures. When subjected 
to a low temperature and high pressure it is converted into 
a liquid; and when some of the liquid is exposed to the air, 
evaporation takes place so rapidly that a great deal of heat 



172 INTRODUCTION TO CHEMISTRY. 

is absorbed, and some of the liquid becomes solid. The 
gas has a slightly acid taste and smell. 

Carbon dioxide is much heavier than air, its specific 
gravity being 1.529. A litre of the gas under standard 
conditions of temperature and pressure weighs 1.977 grams. 
It dissolves in water, one volume of water dissolving about 
one volume of the gas at the ordinary temperature. As is 
the case with all gases, when the pressure is increased the 
water dissolves more gas; and when the pressure is removed 
the gas escapes again. The so-called " soda-water " is 
sinrply water charged with carbon dioxide under pressure. 
The escape of the gas, when the water is drawn, is familiar 
to every one. The carbon dioxide used in charging the 
water is generally made from a sodium salt of carbonic 
acid known as " bicarbonate of soda." 

Chemical Properties of Carbon Dioxide. — Carbon dioxide 
is not combustible, nor does it support combustion. It is 
not combustible for the same reason that water is not; 
because it already holds in combination all the oxygen it 
has the power to combine with. Before it can burn again, 
it must first be decomposed. Carbon has the power to 
combine with oxygen, and in so doing it gives rise to the 
formation of a definite quantity of heat. A kilogram of 
carbon represents a certain quantity of energy, which we 
can get first in the form of heat and then convert into 
other forms, as electricity, motion, etc. After the kilo- 
gram of carbon has been burned, it no longer represents 
the energy it did in the form of carbon. A body of water 
elevated ten or fifteen feet represents a certain quantity of 
energy which can be obtained by allowing the water to fall 
upon the paddles of a water-wheel connected with the ma- 
chinery of a mill. After the water has fallen, however, it 
no longer has power to do work, or it has no energy. In 
order that it may again do work, it must again be lifted. 
Not only does carbon dioxide not burn, but it does not sup- 



RESPIRATION. i73 

port combustion. Although it contains a large quantity of 
oxygen in combination, it does not as a rule give it up to 
other substances. 

[What gas containing oxygen in combination with another 
element does support combustion?] 

Respiration. — It was stated above that carbon dioxide is 
given off from the lungs just as it is from a fire, and the 
fact was demonstrated by means of a simple experiment. 
It is a waste-product of the processes going on in the ani- 
mal body. Just as it cannot support combustion, so also 
it cannot support respiration. It is not poisonous any more 
than water is; but it cannot supply the oxygen which 
is needed for breathing purposes, and hence animals die 
when placed in it. They die by suffocation, as they do in 
drowning. Any considerable increase in the quantity of 
carbon dioxide in the air above that which is normally 
present is objectionable, for the reason that it decreases 
the proportion of oxygen in the air which is breathed. It 
has been found that as much as 5 per cent of pure 
carbon dioxide may be present in air without causing 
injury to those who breathe it. In a badly-ventilated room 
in which a number of people are collected and lights are 
burning, it is well known that in a short time the air 
becomes foul, and bad effects, such as headache, drowsi- 
ness, etc., are produced on the occupants of the room. 
These effects have been shown to be due, not to the carbon 
dioxide, but to other waste-products which are given off 
from the lungs in the process of breathing. The gases 
given off from the lungs consist of nitrogen, oxygen, carbon 
dioxide, and water-vapor. Besides these, however, there 
are many substances in a fine state of division which con- 
tain carbon and are in a state of decomposition. These 
are poisonous, and are the chief cause of the bad effects ex- 
perienced in breathing air which has become contaminated 
by the exhalations from the lungs. As carbon dioxide is 



174 INTRODUCTION TO CHEMISTRY. 

given off from the lungs at the same time, the quantity of 
this gas present is proportional to the quantity of the or- 
ganic impurities. Hence, by determining the quantity of 
carbon dioxide it is possible to determine whether the air 
of a room occupied by human beings is fit for use or not. 
As carbon dioxide is formed in the earth wherever an acid 
solution comes in contact with a carbonate, the gas is 
frequently given off from fissures in the earth. It is hence 
not unfrequently found in old wells which have not been 
in use for some time, and deaths have been caused by 
descending these wells for the purpose of repairing them. 
The gas is also frequently met with in mines, and is called 
choke-damp by the miners. The miners are aware that 
after an explosion caused by fire-damp there is danger of 
death from choke-damp. The reason is simple. When 
fire-damp, or marsh-gas, explodes with air the carbon is 
converted into choke-damp, or carbon dioxide, and the 
hydrogen into water. Air in which a candle will not burn 
is not fit for breathing purposes. 

The Cycle of Carbon in Nature. — The part played by car- 
bon dioxide in nature is extremely important and in- 
teresting. The carbon of living things is obtained from 
carbon dioxide, and returns to this form when life ceases. 
AH living things contain carbon as an essential constitu- 
ent. Whence comes this carbon ? Animals eat either the 
products of plant-life or other animals which derive their 
sustenance from the vegetable kingdom. The food of 
animals comes, then, either directly or indirectly from 
plants. But plants derive their sustenance largely from 
the carbon dioxide of the air. The plants have the power 
to decompose the gas with the aid of the direct light of 
the sun, and they then build up the complex compounds 
of carbon which form their tissues, using for this purpose 
the carbon of the carbon dioxide which they have decom- 
posed. Many of these compounds are fit for food for 



ENERGY STORED IN PLANTS AND ANIMALS, 175 

animals; that is to say, they are of such composition that 
the forces at work in the animal body are capable of trans- 
forming them into animal tissues, or of oxidizing them, and 
thus keeping the temperature of the body up to the neces- 
sary point. That part of the food which undergoes oxida- 
tion in the body plays the same part as fuel in a stove. It 
is burned up with an evolution of heat, the carbon being 
converted into carbon dioxide, which is given off from the 
lungs. From fires and from living animals carbon dioxide 
is returned to the air, where it again serves as food for the 
plants. When the life-process stops in the animal or plant, 
decomposition begins; and the final result of this, under 
ordinary circumstances, is the conversion of the carbon 
into carbon dioxide. 

Plants and Animals as Storehouses of Energy. — Under 
the influence of life and sunlight carbon dioxide is, then, 
converted in the plant into compounds containing carbon 
which are stored up in the plant. These compounds are 
capable of burning, and thus giving heat; or some of them 
may be used as food for animals, assuming other forms 
under the influence of the life-process of the animals. As 
long as life continues, plants and animals are storehouses 
of energy. When death occurs, the carbon compounds 
pass back to the form of carbon dioxide; the energy which 
was stored up is lost. The power to do work which the 
carbon compounds of plants and animals possess comes 
from the heat of the sun. It takes a certain quantity of 
this heat, operating under proper conditions, to decompose 
a certain quantity of carbon dioxide and elaborate the com- 
pounds contained in the plants. When these compounds 
are burned they give out the heat which was absorbed in 
their formation during the growth of the plants. These 
compounds are said to possess chemical energy. This has 
its origin in heat, and is capable of reconversion into heat. 
The transformation of the energy of the sun's heat into 



176 INTRODUCTION TO CHEMISTRY. 

chemical energy lies at the foundation of all life. As the 
heat of the sun acting upon the great bodies of water and 
on the air gives rise to the movements of water which are 
essential to the existence of the world as it is, so the action 
of the sun's rays on carbon dioxide, in the presence of the 
delicate mechanism of the leaf of the plant, gives rise to 
those changes in the forms of combination of the element 
carbon which accompany the wonderful process of life. 

Carbonic Acid and Carbonates. — A solution made by pass- 
ing carbon dioxide into water has a slightly acid reaction. 
[Try it.] It will act upon solutions of bases and form 
salts. The formula of the sodium salt formed in this way 
has been shown to be Na 2 C0 3 ; that of the potassium salt, 
K 2 C0 3 , etc. These salts are plainly derived from an acid, 
H 2 C0 3 , which is carbonic acid. It is probable that this 
acid is contained in the solution of carbon dioxide in water. 
It is, however, so unstable that it readily breaks up into 
carbon dioxide and water : 

H 2 C0 3 t* H 2 + C0 2 . 

When carbon dioxide acts upon a base it forms a salt. 
Thus, with potassium hydroxide or calcium hydroxide the 
action which takes place is represented thus : 

2KOH + CO, = K 2 C0 3 + H 2 0; 

Ca(OH) a + C0 2 =* CaC0 3 + H 2 0. 

With the acid the action would take place as represented 
thus : 

2KOH + H 2 C0 3 m K 2 C0 3 + 2H 2 0; 

Ca(OH), + H 2 C0 3 = CaCOs + 2H 2 0. 

Experiment 94. — Pass carbon dioxide into a solution of caustic 
potash until it will absorb no more. Add acid to some of this 
solution and convince yourself that the gas given off is carbon 



CARBONATES-CARBON MONOXIDE. \11 

dioxide. Write the equations representing the reactions which 
take place on passing the carbon dioxide into the caustic-potash 
solution, and on adding an acid to the solution. What evidence 
have you that the gas given off is carbon dioxide ? 

Experiment 95. — Pass carbon dioxide into 50 to 100 cc. clear 
lime-water. Filter off the white insoluble substance. Try the 
action of a little acid on it. What evidence have you that it is 
calcium carbonate ? How could you easily distinguish between 
lime-water and a solution of caustic potash ? 

Solution of Calcium Carbonate in Water containing Car- 
bon Dioxide. — Although when carbon dioxide is passed into 
lime-water calcium carbonate is precipitated, the calcium 
carbonate dissolves, and the solution finally becomes clear, 
if the gas is passed through it for some time. Water alone 
does not dissolve calcium carbonate, but water containing 
carbon dioxide does. If this solution is heated, the car- 
bon dioxide is driven off and the calcium carbonate is again 
thrown dow T n. Natural waters which flow over limestone 
take up more or less calcium carbonate by virtue of the 
carbon dioxide which they absorb from the air. Such 
waters, w T hich are called hard waters, are in the condition 
of the solution of calcium carbonate above referred to. 
When heated, the calcium carbonate is deposited. This 
is frequently noticed in the deposits in boilers and other 
vessels in which hard water is boiled. 

Experiment 96. — Pass carbon dioxide first through a little 
water to wash it, and then into 50 to 100 cc. clear lime-water. 
At first the insoluble carbonate will come down, as in Experiment 
95 ; but soon it will begin to dissolve, and finally an almost clear 
solution will be obtained. Heat this solution, and the insoluble 
carbonate will again appear. 

Carbon Monoxide, CO. — When a substance containing 
carbon burns in an insufficient supply of air, — as, for ex- 
ample, w^hen the draught in a furnace is not strong enough 
to remove the products of combustion and supply fresh air, 
— the oxidation of the carbon is not complete, and the 



178 INTRODUCTION TO CEEMISTKY. 

product, instead of being carbon dioxide, is carbon mon- 
oxide, CO. This substance can also be made by extracting 
oxygen from carbon dioxide. It is only necessary to pass 
the dioxide over heated carbon, when reaction takes place 
as represented thus : 

C0 3 + C = 2CO. 

This method of formation is illustrated in coal-fires, and 
can be well observed in an open grate. The air has free 
access to the coal, and at the surface complete oxidation 
takes place. But that part of the carbon dioxide which is 
formed at the lower part of the grate is drawn up through 
the heated coal and is partly reduced to carbon monoxide. 
When the monoxide escapes from the upper part of the 
grate it again combines with oxygen, or burns, giving rise 
to the characteristic blue flame always noticed above a 
mass of burning coal. Should anything occur to prevent 
free access of air, carbon monoxide may easily escape com- 
plete oxidation. 

It is also formed by passing water over highly heated 
carbon, when this reaction takes place : 

C + H 2 = CO + 2H. 

This is the reaction which is made use of in the manu- 
facture of " water-gas/' The gas thus obtained is a mix- 
ture of hydrogen and carbon monoxide. Before use it is 
enriched by the addition of hydrocarbons from petroleum. 

Preparation of Carbon Monoxide. — The easiest way to 
make carbon monoxide is to heat oxalic acid, C 2 H 2 4 , with 
five to six times its weight of concentrated sulphuric acid. 
The change is represented thus : 

C 2 H 2 4 = C0 2 + CO + H 2 0. 

Both carbon dioxide and monoxide are formed. Both 
are gases. In order to separate them the mixture is passed 



PROPERTIES OF CARBON MONOXIDE. 179 

through a solution of caustic soda, which takes up the 
carbon dioxide [forming what ?] and allows the monoxide 
to pass. 

Experiment 97. — Put 10 grams crystallized oxalic acid and 50- 
60 grains concentrated sulphuric acid in an appropriate-sized 
flask. Connect with two Wolff's flasks containing caustic-soda 
solution. Heat the contents of the flask gently. Collect some of 
the gas over water. Set fire to some, and notice the characteristic 
blue flame. If convenient put a live mouse in a vessel containing 
a mixture of about equal parts of carbon monoxide and air. It 
will die unless removed. 

Properties of Carbon Monoxide. — Carbon monoxide is a 
colorless, tasteless, inodorous gas, insoluble in water. It 
burns with a pale blue flame, forming carbon dioxide. It 
is exceedingly poisonous when inhaled. Hence it is very 
important that it should not be allowed to escape into 
rooms occupied by human beings. Death is sometimes 
caused by the gases from coal stoves. The most dangerous 
of these gases is carbon monoxide. A pan of smouldering 
charcoal gives off this gas, and the poisonous character of 
the gas is well known, as it has been used to some extent 
for the purpose of suicide. The poisonous character 
of carbon monoxide has led to a great deal of discussion 
and to some legislation on the subject of " water-gas." 
The question has been repeatedly raised whether the manu- 
facture of the gas should be permitted. There is no doubt 
of the fact that it is a dangerous substance ; and that it 
should not be allowed to escape into the air of rooms is 
obvious. With proper precautions, however, there seems 
to be no good reason why it should not be used, although 
it is somewhat more poisonous than coal-gas. 

At high temperatures carbon monoxide combines readily 
with oxvgen, and is hence a good reducing agent. In 
the reduction of iron from its ores, the carbon monoxide 
formed in the blast-furnace plays an important part in the 
reducing process. 



180 INTRODUCTION TO CHEMISTRY. 

Experiment 98.— Pass carbon monoxide over some heated cop- 
per oxide contained in a hard-glass tube. Is the oxide reduced ? 
How do you know ? Is carbon dioxide formed ? What evidence 
have you ? Was the carbon monoxide used free of carbon diox- 
ide? If not, what evidence have you that carbon dioxide is 
formed in this experiment ? 

Illumination, Flame, Blowpipe, etc. — As the substances 
used for illumination contain carbon, and the chemical 
processes involved consist largely in the oxidation of the 
carbon of these compounds, this is an appropriate place to 
take up the subject of illumination, and also that of flame, 
and the blowpipe. 

In all ordinary kinds of illumination flames are the im- 
mediate source of the light. Whether illuminating-gas, a 
lamp, or a candle is used, the light comes from a flame. 
In the first case, the gas is burned directly; in the case of 
the lamp, the oil is first drawn up the wick, then converted 
into a gas, and this burns; while, finally, in the case of the 
candle, the solid material of the candle is first melted, then 
drawn up the wick, converted into gas, and the gas burns, 
forming the flame. In each case, then, there is a burning 
gas, and this burning gas we call a flame. 

Illuminating-gas. — Two kinds of illuminating-gas are 
now made, water-gas and coal-gas. The method of prepa- 
ration and the composition of the former have been given. 
The latter is made from coal by heating in closed retorts. 
When soft coal is heated the hydrogen passes off, partly in 
combination with carbon, as hydrocarbons, and partly in 
the free state. The nitrogen passes off as ammonia, and a 
large percentage of the carbon remains behind in the retort 
in the uncombined state as coke (see page 160). The gases 
given off are purified, and form ordinary illuminating-gas. 
One ton of coal yields on an average 10,000 cubic feet of 
gas. The value of a gas depends upon the amount of light 
given by the burning of a definite quantity. It is meas- 



ILL UMW AUNG GAS-FLAMES. 181 

ured by comparing it with the light given by a candle 
burning at a certain rate. The standard candle is one made 
of spermaceti, which burns at the rate of 120 grains an 
hour. The standard burner used for the gas is one through 
which 5 cubic feet of gas pass in an hour. Now, if it is de- 
sired to determine the illuminating-power of a gas, the gas 
is passed through the standard burner at the rate men- 
tioned, and the light which it gives is compared with the 
light given by the standard candle. The comparison is 
made by means of a so-called photometer. The illuminating- 
power of the gas is then stated in terms of candles. The 
statement that the illuminating-power of a gas is fourteen 
candles means that, when burning at the rate of 5 cubic 
feet an hour, its flame gives fourteen times as much light 
as the standard candle. 

Flames. — Ordinarily when we speak of a flame we mean 
a gas which is combining with oxygen. The hydrogen 
flame is simply the phenomenon accompanying the act of 
combination of the two gases hydrogen and oxygen. Owing 
to the fact that we are surrounded by oxygen, we speak of 
hydrogen as the burning gas. How would it be if we were 
surrounded by an atmosphere of hydrogen? Plainly, oxy- 
gen would then be a burning gas. If a jet of oxygen is 
allowed to escape into a vessel containing hydrogen, a 
flame will appear where the oxygen escapes from the jet, 
if a light is applied. This is an experiment requiring 
great precautions, and, as the principle can be illustrated 
as well by means of illuminating-gas, we may use this in- 
stead. Just as illuminating-gas burns in an atmosphere of 
oxygen, so oxygen burns in an atmosphere of illuminating 



Experiment 99. — Break off the neck of a good-sized retort; fit 
a perforated cork to the small end; pass a piece of glass tube 
through the cork and connect by means of rubber hose with an 
outlet for gas. Fix the apparatus in position, as shown, in Fig. 44, 



182 



INTRODUCTION TO CHEMISTRY. 



Turn the gas on, and when the air is driven out of the retort- 
neck, light the gas. The neck is now filled with illuminating-gas, 
and the gas is burning at the mouth of the vessel. If now a 
platinum jet from which oxygen is issuing is passed up into the 




Fig. 44. 

gas the oxygen will take fire, and a flame will appear where the 
oxygen escapes from the jet. The oxygen burns in the atmos- 
phere of coal-gas. 

Kindling-temperature of Gases. — In studying the action 
of oxygen upon other substances, we learned that it is neces- 
sary that each of these substances should be raised to a 
certain temperature before it will combine with the oxygen. 
This statement is as true of gases as of other substances. 
When a current of hydrogen is allowed to escape into the 
air, or into oxygen, no action takes place unless it is heated 
up to its burning-temperature, when it takes fire and con- 
tinues to burn, as the burning of one part of the gas heats 
up the part which follows it, and hence the gas is heated 
up to the burning-temperature as fast as it escapes into the 
air. If the gas should be cooled down even very slightly 
below this temperature, it would be extinguished. This is 
shown in a very striking manner by the following experi- 
ments : 



SAFETY-LAMP. 



183 



Experiment 100. — Light a Bunsen burner. Bring down upon 
the flame a piece of fine brass or iron wire gauze. There is no 
flame above the gauze. That the gas passes through unburned 
can be shown by applying a light just above the outlet of the 
burner and above the gauze. The gas will take fire and burn. 
By simply passing through the thin wire gauze, then, the gas is 
cooled down below its burning-temperature, and does not burn 
unless it is heated up again. Turn on a Bunsen burner. Do not 
light the gas. Hold a piece of wire gauze about one and a half to 
two inches above the outlet. Apply a lighted match above the 
gauze, when the gas will burn above the gauze, but not below it. 
Here again the heat necessary to raise the temperature of the gas 
to the burning-temperature cannot be communicated through the 
gauze. If in either of the above-described experiments the gauze 
is held in position for a time, it will probably become so highly 
heated that the gas on the one side where there is no flame will 
be raised to the burning-temperature. The instant that point is 
reached the flame becomes continuous. 

Safety-lamp. — The principle illustrated in the preceding 
experiments is utilized in the miner's safety -lamp. One of 
the dangers which the coal-miner has to encounter is the 
occurrence of fire-damp, or methane, CH 4 , 
which with air forms an explosive mixture. 
The explosion can be brought about only 
by contact of flame with the mixture. In 
order to avoid the contact, the flame of 
the safety-lamp is surrounded by wire 
gauze, as shown in Fig. 45. When a lamp 
of this kind is brought into an explosive 
mixture of marsh-gas and air, the mixture 
passes through the wire gauze and comes 
in contact with the flame, and a slight ex- 
plosion occurs inside the gauze, but the 
flame of the burning gas inside the wire 
gauze cannot pass through and raise the 
temperature of the gas outside to the burn- 
ing-temperature. Hence no serious ex- 
plosion can take place. The flickering of the flame of the 




Fig. 45. 



184 INTRODUCTION TO CHEMISTRY. 

lamp, and the occurrence of small explosions inside, fur- 
nish the miner with the information that he is in a danger- 
ous atmosphere. 

m 

Structure of Flames.— The hydrogen flame consists of a 
thin envelope of burning hydrogen enclosing unburned 
gas, and surrounded by water-vapor, which is the product 
of the combustion. The structure of other flames depends 
upon the complexity of the gases burned and the condi- 
tions under which the burning takes place. In general, a 
flame consists of an outer envelope of gas combining with 
oxygen, and hence hot, and an inner part which contains 
unburned gas, which is, for the most part, cool. A part of 
the unburned gas is, however, hot, and it would combine 
with oxygen were it not for the fact that it is surrounded 
by an envelope which prevents access of air. The outer 
hot part of the flame is called the oxidizing flame, because 
it presents conditions favorable to the oxidation of sub- 
stances introduced into it. The inner hot part is called 
the reducing flame, because it consists of highly-heated 
substances which have the power to combine with oxygen, 
and hence many compounds containing oxygen 
[ lose it, or are reduced, when introduced into this 
part of the flame. The hottest part of the flame 
is at the extreme top. Here oxidation is taking 
place most energetically. The hottest part of the 
unburned gases is at the tip of the dark central 
j part of the flame. In the flame of a Bunsen 
burner the two parts can be distinguished very 
easily. The dark central part of the flame ex- 
__ B tends for some distance above the outlet of the 
burner. If the holes at the base of the burner 
are partly closed, the tip of the central part of 
the flame becomes luminous. This luminous tip 
fig. 46. j s mos t efficient for the purpose of reduction. 
The principal parts of the flame are those marked in 



BLOWPIPE. 185 

Fig. 46. B is the central cone of unburned gas. C is the 
luminous tip, the best part of the flame for reduction. A 
is the envelope of burning gas. This is further surrounded 
by a non-luminous envelope consisting of the products 
of combustion, carbon dioxide and water-vapor. Certain 
metals placed in the upper end of the flame take up oxy- 
gen, because they are highly heated in the presence of 
oxygen. Certain oxides lose their oxygen when placed 
in the tip of the central cone, because the gases are here 
heated to the temperature at which they have the power 
to combine with oxygen. 

Blowpipe. — The oxidizing and reducing flames are fre- 
quently utilized in the laboratory. For the purpose of 
increasing their efficiency a blowpipe is used. This is a 
tube through which air is blown into a flame by means of 
the mouth. It is usually constructed in the shape shown 
in Fig. 47. At the smaller end, which is placed in the 



p^= 



Fig. 47. 

flame, there is usually a small tube of platinum. The 
blowpipe may be used with the flame of a candle, an 
alcohol-lamp, or a gas-lamp. It is most frequently used 
with the gas-lamp. A piece of brass tubing which fits 
snugly in the tube of a Bunsen burner is cut off and ham- 
mered together so as to leave a narrow slit-like opening. 
This tube is then slipped into the burner, as shown in Fig. 
48. It reaches to the bottom of the burner, and thus 
cuts off the supply of air which usually enters the holes at 
the base. The gas is now lighted and the current so regu- 
lated that there is a small flame about 1| to 2 inches high. 
The tip of the blowpipe is placed on the slit of the burner 



186 



INTRODUCTION TO CHEMISTRY. 



in the flame, so that it extends about one third the way 
across it, as shown in Fig. 49. By blowing regularly and 
not too violently through the pipe the flame is forced down 
in the same direction as the end-piece of the blowpipe, and 
the slant of the burner-slit. Under proper 
conditions it separates sharply into a central 
blue part and an outer part of another color. 
The direction and lines of division of the flame 






Fig. 48. 



Fig. 49. 



Fig. 50. 



are indicated in Fig. 50. The extreme outer tip A is the 
most efficient oxidizing flame. The tip B of the inner 
blue part is the most efficient reducing flame. 

The use of the blowpipe is illustrated by the following 
experiments : 

Experiment 101. — Select a piece of charcoal about 4 inches 
long by 1 inch wide and 1 inch thick, with one surface plane.* 
Near the end of the plane surface make a cavity by pressing the 
edge of a small coin against it, and turning it completely round 
a few times. Mix together equal small quantities of dry sodium 
carbonate and lead oxide. Put a little of the mixture in the 
cavity in the charcoal, and heat it in the reducing flame pro- 
duced by the blowpipe. In a short time globules of metallic 
lead will be seen in the molten mass. After cooling, scrape the 
solidified substance out of the cavity in the charcoal. Put it into 
a small mortar, treat it with a little water, and, after breaking 
it up and allowing as much as possible to dissolve, pick out the 



* Pieces of charcoal prepared for blowpipe work can be bought 
from dealers in chemical apparatus at small cost. 



BLOWPIPE— THE LUMINOSITY OF FLAMES. 187 

metallic beads. [Is it malleable or brittle ? Is metallic lead 
malleable or brittle ? Is it dissolved by hydrochloric acid ? Is 
lead soluble in hydrochloric acid? Is it soluble m nitric acid? 
Is lead soluble in nitric acid ?] The action of the acids may be 
tried by putting the bead on a small dry watch-glass and adding 
a few drops of the acid. [Does the substance act like lead? 
What has become of the oxygen with which the lead was com- 
bined in the oxide ? Is there any special advantage in having a 
support of charcoal for this experiment ?] 

Experiment 102. — Heat a small piece of metallic lead on char- 
coal in the oxidizing flame of the blowpipe. Notice the formation 
of the oxide, which forms a coating or film on the charcoal in the 
neighborhood of the metal. [Is there any analogy between this 
process and the burning of hydrogen ? In what does the analogy 
consist ? What differences are there between the two processes ?] 

Use of the Blowpipe in Analysis. — Some oxides are re- 
duced very easily when heated in the reducing flame. 
Others are not. The composition of a substance can 
often be determined by heating it in the blowpipe flame 
and no'ticing its conduct. Some metals are easily oxi- 
dized in the oxidizing flame. Some form characteristic 
films of oxides on the charcoal, and in some cases it is possi- 
ble to detect the presence of certain substances by noticing 
the color of the film of oxide. The blowpipe is therefore 
of great value as affording a means of detecting the presence 
of certain elements in mixtures or compounds of unknown 
composition. The chemical principles involved in its use 
will be clear from what has already been said. 

Causes of the Luminosity of Flames. — It is evident from 
what has been seen that flames vary greatly in their light- 
giving power. The hydrogen flame, for example, gives 
practically no light. This is also the case with the flame 
of the Bunsen burner ; while, on the other hand, the flame 
of illuminating-gas burning under ordinary circumstances, 
and that of a candle, etc., give light. [To what is the 
difference due ?] There are several causes which operate tc 



188 INTRODUCTION TO CHEMISTRY. 

make a flame give light, and vice versa. In the first place, 
if a solid substance which does not burn up is introduced 
into a non-luminous flame, a part of the heat appears as 
light. This is seen when a spiral of platinum wire is in- 
troduced into a hydrogen flame. It has also been shown 
by introducing a piece of lime into the hot non-luminous 
flame of the oxyhydrogen blowpipe. A similar cause 
operates in ordinary gas-flames to make them luminous. 
There are always present particles of unburned carbon, as 
can be shown by putting a piece of porcelain or any solid 
substance into the flame, when there will be deposited on it 
a layer of soot, which consists mainly of finely-divided car- 
bon. In the flame these particles of carbon are heated to 
the temperature at which they give light. Again, it has 
been found that the same candle gives more light at the 
level of the sea than it does when at the top of a high moun- 
tain, as Mont Blanc, on which the experiment was actually 
performed. This is partly due to a difference in the density 
of the gases. Naturally, the denser the gas the more ac- 
tive the combustion, the greater the heat, and the greater 
the light. This last statement ceases to be true when the 
oxidation becomes sufficient to burn up all the solid parti- 
cles of carbon in the flame. If gases which in burning 
give light are cooled down before they are burned, the 
luminosity is diminished, and, conversely, non-luminous 
flames may be rendered luminous by heating the gases be- 
fore burning them. When gases w T hich give luminous 
flames are diluted to a sufficient extent with neutral gases, 
such as nitrogen and carbon dioxide, which neither burn 
nor support combustion, they become non-luminous. 

Bunsen Burner. — All the statements made in regard to 
the causes of the luminosity of flames are based upon care- 
fully-performed experiments. These experiments, how- 
ever, cannot be easily repeated by the student in the lab- 
oratory in a satisfactory way. One constant reminder of 



BUNSEN BURNER-CYANOGEN. 189 

the possibility of rendering a luminous flame non4uminoug 5 
and vice versa, is furnished by the burner universally used 
in chemical laboratories, and called, after the name of its 
inventor, the Bunsen Burner. The construction of this 
burner is easily understood. It consists of a base and an 
upper tube. The base is connected by means of a rubber 
tube with the gas supply. The gas escapes from a small 
opening in the base, and passes up through the tube. At 
the lower part of the tube there are two holes, which may 
be opened or closed by turning a ring with two correspond- 
ing holes in it. When the gas is turned on, it is lighted at 
the top of the tube. Air is at the same time drawn 
through the holes at the base. The result is that the 
flame is practically non-luminous. If the ring at the base 
is turned so that the air-holes are closed, the flame be- 
comes luminous. The advantage of the non-luminous 
flame for laboratory use consists in the fact that it does 
not deposit soot, and, at the same time, gives a good heat. 

[Could the hydrogen flame deposit soot ?] 

The non-luminosity of the flame of the Bunsen burner 
appears to be due to several causes : (1) Dilution of the 
gases by means of the nitrogen of the air ; (2) Cooling of 
the gases by the entrance of the air; (3) Burning of the 
solid particles by the aid of the oxygen of the air admitted 
to the interior of the flame. 

Cyanogen, C 4 N 2 . — Carbon does not combine with nitro- 
gen under ordinary circumstances. If, however, they are 
brought together at very high temperatures in the pres- 
ence of metals, they combine to form compounds known 
as cyanides. Thus, when nitrogen is passed over a highly- 
heated mixture of carbon and potassium carbonate, K 2 C0 3 , 
the compound potassium cyanide, KCN, is formed. Car- 
bon containing nitrogen, as animal charcoal, when ignited 
with potassium carbonate, reduces the potassium carbonate, 
forming potassium, and this causes the carbon and nitrogen 



190 introduction to chemistry. 

to combine, forming potassium cyanide. When refuse 
animal substances, such as blood, horns, claws, hair, wool, 
etc., are heated together with potassium carbonate and 
iron, a substance known as potassium ferrocyanide, or 
yelloio prussiate of potash, 4KCN.Fe(CN) 2 + 3H 2 0, is 
formed. When this is simply heated it decomposes, 
yielding potassium cyanide. It is not a difficult matter 
to make mercuric cyanide, Hg(CN) 2 , from the potassium 
compound. By heating mercuric cyanide it breaks up 
yielding metallic mercury and cyanogen gas: 

Hg(CN) 2 = Hg + C,N 2 . 

[What analogy is there between this reaction and that 
which takes place when mercuric oxide is heated ?] 

Cyanogen is a colorless gas. It receives its name from 
the fact that some of its compounds are blue (kvcxvos, 
blue). It is easily soluble in water and alcohol. It is 
extremely poisonous. 

Hydrocyanic Acid, Prussic Acid, HCN. — This acid oc- 
curs in nature in combination with other substances, — in 
bitter almonds, the leaves of the cherry, laurel, etc. It is 
prepared by treating the cyanides with sulphuric or hy- 
drochloric acid. Thus, by treating potassium cyanide with 
sulphuric acid this reaction takes place : 

2KCN + H 2 S0 4 = K 2 S0 4 + 2HCN. 

[What reactions already considered does this suggest ?] 
Further, by treating potassium cyanide with a solution 
of hydrochloric acid in water, hydrocyanic acid is liberated : 

KCN + HC1 = KC1 + HCN. 

[Compare these reactions with similar reactions already 
studied.] 



SUMMARY. 191 

Hydrocyanic acid is a volatile liquid which boils at 26.5°, 
and solidifies at — 15°. It has a very characteristic odor, 
resembling that of bitter almonds. It is extremely poi- 
sonous. It dissolves in water in all proportions, and it is 
such a solution that is known as prussic acid. 

Both cyanogen and hydrocyanic acid are extremely un- 
stable. In the presence of water, the nitrogen tends to 
combine with hydrogen to form ammonia, and the carbon 
with oxygen and hydrogen to form more stable compounds. 

Summary. — Carbon is contained in all living things, and 
in their fossil remains. The number of compounds which 
it forms is almost infinite. They are usually treated to- 
gether under the head of Organic Chemistry. 

Carbon is found in the atmosphere in the form of car- 
bon dioxide, and in the form of carbonates is widely dis- 
tributed in the earth. 

Uncombined, it occurs in nature as diamond and 
graphite. 

Amorphous carbon is a third variety of carbon. Char- 
coal in its various forms is amorphous carbon. It is made 
by charring organic substances which contain carbon, hy- 
drogen, and oxygen. Coke, lamp-black, and bone-black 
are other forms of amorphous carbon. Bone-black has the 
power to extract coloring matters from solutions. Char- 
coal has the power to absorb gases, and is used for purify- 
ing air. It also absorbs disagreeable substances from 
water, and is used for the purpose of purifying water. 

Coal is a form of carbon found in nature in many varie- 
ties. The soft coals contain more hydrogen than the hard 
coals, which contain a larger percentage of carbon. 

At ordinary temperatures carbon is a very inactive ele- 
ment. At high temperatures it combines with oxygen 
with avidity. It is hence a good reducing agent, and is 
used extensively as such in the extraction of metals from 
their ores. 



192 INTRODUCTION TO CHEMISTRY. 

Carbon forms a large number of compounds with hydro- 
gen. These are the hydrocarbons. 

Carbon dioxide is formed in many natural processes, as 
in respiration, combustion, decay, and fermentation. It is 
prepared by treating a carbonate with an acid. The gas 
given off is not carbonic acid, but a substance which bears 
to the acid the relation of an anhydride. 

Carbon dioxide is the food of plants. Plants form the 
food of animals. Animals give back carbon dioxide to the 
air in the process of breathing. After death the carbon of 
animals and plants, if left exposed to the air, passes back 
to the form of carbon dioxide, and again starts on its 
round. 

Carbon dioxide forms salts, with bases. These have the 
general formula M 2 C0 3 , in which M represents any metal, 
such as potassium, sodium, etc. These are very unstable, 
being decomposed by any acid. 

Calcium carbonate is insoluble in water, but it dissolves 
in water containing carbon dioxide. When heated the 
carbon dioxide is driven off and the calcium carbonate 
deposited. This phenomenon is the same as that which 
gives rise to the ordinary boiler incrustations. 

Carbon monoxide is a poisonous gas, which is formed by 
incomplete oxidation of carbon or incomplete reduction of 
carbon dioxide. It is formed in ordinary coal fires by the 
passage of carbon dioxide over thoroughly heated coal. It 
combines readily with oxygen, and is hence a good reduc- 
ing agent. 

A flame is a burning gas. A gas that burns in oxygen 
will form an atmosphere in which oxygen will burn. If a 
burning gas is cooled down even very slightly below its 
burning-temperature, it is extinguished. In the miner's 
safety-lamp the flame is surrounded by a piece of wire 
gauze. The gas cannot pass through this gauze without 
being cooled down below the burning-temperature. 

Flames are made up of different parts with different 



8 VMM ART. 193 

properties. The outer tip is the hottest part, and is called 
the oxidizing flame. The tip of the dark inner part, con- 
sisting of unburned gas, is the reducing flame. 

A luminous flame can be made non-luminous by diluting 
the burning gas with neutral gases ; by cooling the gases ; 
by introducing oxygen into the gas so as to effect complete 
oxidation of the carbon. 

In the presence of metals carbon and nitrogen combine 
to form cyanides. From these, cyanogen and hydrocyanic 
acid are obtained. 



CHAPTER XIII. 

AVOGADRO'S LAW.— MOLECULAR WEIGHTS.-MOLECULAR 
FORM ULAS. —VALENCE . 

Avogadro's Law. — Early in this century the Italian 
physicist and chemist, Avogadro, occupied himself with 
the study of the specific gravities of gaseous substances, 
and he recognized clearly that there is some connection 
between the figures representing the relative weights of 
equal volumes of gases and those representing the combin- 
ing weights. It has already been pointed out that the 
weights of equal volumes of hydrogen, chlorine, and oxy- 
gen bear to one another the same relation as their atomic 
weights, viz., 1:35.4:16. The same relation is noticed in 
the case of other gases. This fact, taken together with 
others relating to the physical properties of gases, led Avo- 
gadro to the conception that equal volumes of all gases 
under the same conditions of temperature and pressure 
contain the same number of molecules, the molecule of a 
substance being the smallest particle of that substance as 
it exists in the free state or uncombined. This is known 
as Avogadro's law. It has been tested in a great many 
ways, and has always asserted itself as correct. The inves- 
tigations of both chemists and physicists have only tended 
to confirm it, and at the present day it forms one of the 

194 



AVOGRADVS LAW. 195 

most important foundations of thought in regard to chem- 
ical phenomena. 

The weights of equal volumes of gases can be determined 
without difficulty. According to the law of Avogadro, 
these weights bear to one another the same relation that 
the weights of the molecules of these substances do. Take, 
for example, those compounds thus far studied which are 
gases at ordinary temperatures, or can easily be converted 
into gases by heat. These are water, hydrochloric acid, am- 
monia, nitrous oxide, nitric oxide, marsh-gas, carbon dioxide, 
carbon monoxide, cyanogen, hydrocyanic acid. The specific 
gravities of these substances in the form of gas or vapor 
have been determined. They are: water, 0.623; hydro- 
chloric acid, 1.247; ammonia, 0.597; nitrous oxide, 1.520; 
nitric oxide, 1.039; marsh-gas, 0.557; carbon dioxide, 
1.529; carbon monoxide, 0.968; cyanogen, 1.8; hydrocy- 
anic acid, 0.918. These figures express the relative weights 
of equal volumes of the gases, and they also express the 
relation between the weights of the molecules of the sub- 
stances. It is only necessary to adopt some standard to 
which we can refer the weights of other molecules. Hydro- 
chloric acid will serve the purpose. The smallest molecular 
weight that can be assigned to this compound without 
making the atomic weight of hydrogen less than unity is 
36.1, for hydrochloric acid consists of 1 part by weight of 
hydrogen combined with 35.4 parts by Weight of chlorine. 
Hence, if the sum of the weights of its atoms or its molec- 
ular weight were less than 36.4, the weight of the atom of 
hydrogen would be less than 1. If the molecular weight 
of hydrochloric acid is 36.4, it is an easy matter to calcu- 
late the molecular weights of the other substances men- 
tioned, for, according to Avogadro's hypothesis, they bear 
to the molecular weight of hydrochloric acid the same 
relation that their specific gravities bear to the specific 
gravity of hydrochloric acid. The results of the calcula- 
tion are given in the subjoined table : 



196 



INTRODUCTION TO CBEMiSTMT. 



Compound. 



Water 

Hydrochloric acid 
Ammonia. ....... 

Nitrous oxide. . . . 

Nitric oxide 

Marsh-gas 

Carbon dioxide. . . 
Carbon monoxide. 

Cyanogen 

Hydrocyanic acid 



Sp. Gr. of Gas 


Calculated 


or Vapor. 


Molec. Weight. 


0.623 


18.1 


1.247 


36.4 


0.597 


17.4 


1.52 


44.2 


1.039 


30.2 


557 


16.2 


1.529 


44.4 


0.968 


28.2 


1.8 


52.4 


0.948 


27.6 



The figures thus obtained are relatively correct, provided 
always the law upon which the calculation is based is cor- 
rect. Now, by analysis, it is found that in 18 parts by 
weight of water there are 2 of hydrogen and 16 of oxygen; 
in 36.4 parts by weight of hydrochloric acid there are 1 of 
hydrogen and 35.4 of chlorine; in 17 of ammonia, 14 of 
nitrogen and 3 of hydrogen; in nitrous oxide, 28 of nitro- 
gen and 16 of oxygen; in nitric oxide, 14 of nitrogen and 
16 of oxygen; in marsh-gas, 12 of carbon and 4 of hydrogen; 
in carbon dioxide, 12 of carbon and 32 of oxygen; in car- 
bon monoxide, 12 of carbon and 16 of oxygen; in cyano- 
gen, 24 of carbon and 28 of nitrogen; in hydrocyanic 
acid, 1 of hydrogen, 12 of carbon, and 14 of nitrogen. 
Knowing the weights of the molecules into which an ele- 
ment enters, and the relative weight of the element present 
in these molecules, the smallest weight of the element that 
enters into the composition of molecules is selected as the 
atomic weight. Thus, the examination of all known oxy- 
gen compounds that can be studied in the form of gas or 
vapor shows that the smallest weight of oxygen found in 
any molecule is represented by 16, using the standard 
already adopted. Thus, in water, to make up the molecu- 
lar weight, 18, there are 16 parts by weight of oxygen and 
2 parts by weight of hydrogen; in nitrous oxide, 28 of ni- 



MOLECULES OF THE ELEMENTS. 197 

trogen and 16 of oxygen; in carbon dioxide, 12 of carbon 
and 32 of oxygen; in carbon monoxide, 12 of carbon and 
16 of oxygen. The number 16 is therefore selected as the 
atomic weight of oxygen. 

The ratio of the specific gravity of a gas to its molecular 
weight is approximately 1 : 28.88, i.e., 

~ =b 28.88, or M=dX 28.88, 
ct 

in which M represents the molecular weight of a gaseous 
compound, and d its specific gravity as compared with air 
as the standard. This gives the molecular weight very 
nearly. The exact figure to be adopted is then determined 
by analysis. 

Molecules of the Elements. — The acceptance of Avoga- 
dro's law leads to a curious conclusion regarding the struc- 
ture of elementary gases. The molecular weights of hy- 
drogen, oxygen, chlorine, and nitrogen are found to be 
2, 32, 70.8, and 28 respectively. In other words, they are 
twice as great as their atomic weights. According to this, 
these gases consist of molecules which are twice as heavy as 
their atoms, or, in other words, the molecules of these 
elementary gases consist of two acorns each. The same 
conclusion is reached by another line of reasoning. When 
one volume of hydrogen combines with one volume, two 
volumes of hydrochloric acid are formed. Now, as equal 
volumes of all gases contain the same number of molecules, 
if we assume that in a certain volume of hydrogen there 
are 100 molecules, then in the same volume of chlorine 
and of hydrochloric acid there are also 100 molecules. But 
from 1 volume containing 100 molecules of hydrogen and 
1 volume containing 100 molecules of chlorine 2 volumes 
containing 200 molecules of hydrochloric acid are formed. 
In each molecule of hydrochloric acid gas there must 
be at least one atom of hydrogen and one atom of 



198 INTRODUCTION TO CHEMISTRY. 

chlorine, and in the 200 molecules of hydrochloric acid 
there must be 200 atoms of hydrogen and 200 atoms of 
chlorine. These 200 atoms of hydrogen, however, must 
have been contained in the 100 molecules of hydrogen with 
which we started, and similarly the 200 atoms of chlorine 
must have been contained in the 100 molecules of chlorine. 
Therefore, each molecule of hydrogen must consist of at 
least 2 atoms of hydrogen, and each molecule of chlorine 
must consist of at least 2 atoms of chlorine. 

A similar study of other elementary gases leads to similar 
conclusions in regard to their molecules. The molecule of a 
few elementary gases has been shown to consist of 4 atoms, 
that of some of 3 atoms,* and of a few others of a single 
atom; but usually the condition is that found in hydrogen 
and chlorine. The view is thus forced upon us that the 
molecules of elementary gases consist of atoms of the same 
kind, just as the molecules of compound gases consist of 
atoms of different kinds. The molecule of hydrogen is a 
compound of two atoms of hydrogen, just as the molecule 
of hydrochloric acid is a compound of an atom of hydrogen 
and an atom of chlorine. According to this conception, 
when hydrogen gas and chlorine gas are brought together, 
the complete action is not represented by the equation 

H + C1 = HC1. 



* In speaking of ozone, it was stated that when oxygen is changed 
to ozone there is a diminution of volume from three to two without 
change of weight. In other words, the specific gravity of oxygen 
is two thirds that of ozone. But the specific gravity of oxygen leads 
to the conclusion that its molecule contains two atoms. Similarly, 
the specific gravity of ozone leads to the conclusion that its molecule 
contains three atoms. Ozone is therefore believed to be made up 
of molecules each of which consists of three atoms of oxygen ; and 
ordinary oxygen to be made up of molecules each of which consists 
of two atoms of oxygen. The molecular weight of ordinary oxygen 
is 32 3 and that of ozone is 48, 



MO LEG ULES-NASCENT ST A TE. 199 

The molecules of hydrogen and chlorine must be broken 
up before the act of combination can take place. Hence, 
there are two acts involved in passing from hydrogen gas 
and chlorine gas to hydrochloric acid. These are : 

HH + C1C1 = H + H + CI + CI. 

Molecule of Molecule of Atoms of 

hydrogen. chlorine. hydrogen. 

Then, further, the atoms combine to form compound mole- 
cules : 

H + H + CI + CI = 2HC1. 
Or we may write the equation thus : 

H 2 + Cl 2 = 2HC1. 

Molecule of Molecule of 

hydrogen. chlorine. 

Again, when an elementary gas such as hydrogen or oxy- 
gen is set free from a compound, it appears from the above 
that, at the instant it is liberated, it exists in the atomic 
condition, but that if there is nothing else present with 
which the atoms can combine, they combine with each 
other to form molecules. After it has been set free, there- 
fore, it should be less active than at the instant it is set 
free. This is quite in accordance with many curious and 
well-known facts. 

Nascent State. — It is found that at the instant elements 
are set free from their compounds they are capable of 
effecting changes which they cannot, effect after they have 
once been set free. Thus, free oxygen gas passed into hy- 
drochloric acid produces no change under ordinary condi- 
tions ; but oxygen liberated from a compound in contact 
with hydrochloric acid decomposes the latter and sets 
chlorine free. Hydrogen gas passed into nitric acid causes 
no change ; but hydrogen liberated m direct contact with 



200 INTRODUCTION TO CHEMISTRY. 

nitric acid reduces the acid to ammonia in some cases. 
Many other examples of this kind of action might be cited. 
The simplest explanation of the phenomenon is that offered 
above. An element at the instant of its liberation is said 
to be in the nascent state. 

Relation of Physics and Chemistry to Molecules. — Ac- 
cording to what has been said, all substances, elementary 
as well as combined, are made up of molecules. The 
molecules are believed to have the properties of the sub- 
stance as we know it in the free state. The molecule 
is the smallest particle of a substance that can exist in 
the free state. The molecules are said to be held together 
by cohesion, and, theoretically, a substance could be sepa- 
rated into its molecules by purely mechanical processes. 
As long as action upon a substance does not involve de- 
composition of the molecules, the action is in the realm 
of physics. The molecules are made up of atoms. The 
atom enters into chemical action and is the smallest par- 
ticle of a substance that can do so. Chemistry is that 
science which has to deal with changes within the mole- 
cules. It must be remembered that these statements are 
not statements of facts known to us. The laws of definite 
and multiple proportions are statements of facts ; but when 
we come to speak of atoms and molecules we are dealing 
with conceptions which, however probable they may ap- 
pear, can nevertheless not be proved to be true. We make 
use of these conceptions because they simplify our dealings 
with the facts of chemistry, and suggest lines of inquiry 
which lead to discoveries of value. 

Explanation of the Laws Governing the Combination of 
Gases.— It has been pointed out (pages 140-141) that when 
hydrogen combines with chlorine, with oxygen, and with 
nitrogen the relations between the volumes of the combin- 
ing gases are simple, and that these volumes in turn bear a 
simple relation to the volumes of the products formed, 



HOW A FORMULA IS DETERMINED. 201 

The explanation of these facts on the basis of Avogadro's 
law is as follows : 

In equal volumes of hydrogen and of chlorine there is 
the same number of molecules. Each molecule of hydro- 
gen and each molecule of chlorine consists of two atoms. 
When hydrogen and chlorine combine, one atom of one 
combines with one of the other, so that from one molecule 
of hydrogen and one of chlorine two molecules of hydro- 
chloric-acid gas are formed. The number of molecules is 
the same after combination as before, and therefore the 
product occupies the same volume as the uncombined 
gases. When hydrogen combines with oxygen, however, 
two atoms of hydrogen combine with one of oxygen. The 
reaction is represented thus: 

2H 2 + 2 = 2H 2 0. 

In this case two molecules of hydrogen and one molecule 
of oxygen give two molecules of water, and the volume of 
the product in the form of vapor is only two thirds that of 
the combining gases. The reaction between nitrogen and 
hydrogen is represented thus : 

3H, + N 2 = 2NH 3 . 

Or, from four molecules only two are obtained. Conse- 
quently the volume of the product is only half that of the 
uncombined gases. 

How a Formula is Determined. — Chemical formulas were 
first introduced for the purpose of expressing the composi- 
tion of substances. They might be used for this purpose 
at present without having any connection whatever with 
the conception of atoms and molecules, but the difficulty 
would then be to decide upon the combining weights of 
the elements. It would be possible for authoritative bodies 
to unite in issuing an edict that the combining weights of 
the elements shall be certain figures which are in harmony 



202 INTRODUCTION TO CHEMISTRY. 

with facts known. But this would hardly be a scientific 
mode of procedure; and there might exist differences of 
opinion in regard to the advisability of accepting the fig- 
ures. When, however, we once accept the atomic theory 
and the law of Avogadro, we have a definite basis to work 
on, and there is little opportunity for disagreement in re- 
gard to the figures to be adopted. 

The necessary steps in the determination of the formula 
of a compound may be illustrated by the case of water. 
The compound is first analyzed and found to contain hy- 
drogen and oxygen in the proportion of 1 part hydrogen 
to 8 parts oxygen. This is a fact. But we wish to express 
by our formula not only the composition of the substance, 
but the composition of a molecule of the substance. We 
therefore determine the molecular weight by the method 
described above by comparing the specific gravity of its 
vapor with that of hydrochloric acid or hydrogen. We find 
that the molecular weight is 18. In other words, the mole- 
cule of water, or the smallest particle of water, is 18 times 
heavier than an atom of hydrogen. According to the 
analysis, the 18 parts are made up of 2 parts of hydrogen 
and 16 parts of oxygen. By an examination of a large 
number of gaseous compounds containing oxygen we con- 
clude that 16 is the atomic weight of oxygen, as the 
smallest weight of oxygen found in any of its compounds 
is 16 times heavier than the smallest weight of hydrogen 
found in any of its compounds. Therefore, the molecule 
of water consists of 2 atoms of hydrogen and 1 atom of 
oxygen. The formula representing the facts and concep- 
tions in regard to the composition of water is H 2 0. 

Every formula is intended to express the composition 
and relative weight of a molecule of the compound repre- 
sented. But only in the case of compounds which are 
gases or which can be converted into vapors have we a 
definite basis for assuming that the formulas do represent 
the relative weights of the molecules. 



VALENCE. 203 

Valence.— The formulas of the compounds thus far 
studied have all been determined by exactly the same 
methods. On comparing the formulas of the hydrogen 
compounds of chlorine, oxygen, nitrogen, and carbon, one 
cannot fail to be struck by certain curious differences be- 
tween them. The formulas are 

C1H, OH a , NH al CH 4 . 

Speaking in terms of the theory, the molecule of hydro- 
chloric acid consists of 1 atom of chlorine combined with 1 
atom of hydrogen; the molecule of water consists of 1 
atom of oxygen combined with 2 atoms of hydrogen; the 
molecule of ammonia consists of 1 atom of nitrogen com- 
bined with 3 atoms of hydrogen; the molecule of marsh- 
gas consists of 1 atom of carbon combined with 4 atoms of 
hydrogen. It is thus seen that the atoms of chlorine, 
oxygen, nitrogen, and carbon differ in their power of hold- 
ing hydrogen in combination. The oxygen atom has twice 
the power of the chlorine atom, the nitrogen atom has 
three times this power, and the carbon atom has four times 
this power. An examination of the compounds of other 
elements shows that other atoms differ from one another 
in the same way. 

The smallest power, as far as the number of other atoms 
which it can hold in combination is concerned, is that of 
the chlorine atom. And as one chlorine atom can hold 
but one atom of hydrogen in combination, so one atom of 
hydrogen can hold but one atom of chlorine in combina- 
tion. Either the hydrogen atom or the chlorine atom may 
be taken as an example of the simplest kind of atom. An 
element like hydrogen and chlorine is called a univalent 
element; an element like oxygen whose atom can hold two 
unit atoms in combination is called a bivalent element; an 
element like nitrogen whose atom can hold three unit 
atoms in combination is called a trivalent element; an ele- 



204 INTRODUCTION TO CHEMISTBT. 

ment like carbon whose atom can hold four unit atoms in 
combination is called a quadrivalent element. Most ele- 
ments belong to one or the other of these four classes, 
though there are some which can hold five, six, and even 
seven unit atoms in combination. These are, however, 
rare, and for our present purpose they will require but 
slight notice. 

Valence is that property of an element by virtue of which 
its atom can hold a definite number of other atoms in com- 
bination. ^U^wViOj^wA^ 

[Calcium forms with chlorine the compound CaCl 2 . 
What is the valence of calcium ? Potassium and sodium 
form chlorides of the formulas KC1 and NaCl respectively. 
What is the valence of these elements ? Sulphur forms 
with hydrogen a compound of the formula SH 2 . What is 
the valence of sulphur ?] 

Substituting Power of Elements. — It has been shown 
that, in the formation of salts, the hydrogen of acids is 
replaced by metals. In such cases one atom of a univalent 
metal takes the place of one atom of hydrogen, one atom 
of a bivalent metal takes the place of two atoms of hydrogen, 
etc. Thus, potassium and sodium are univalent. An atom 
of either takes the place of one atom of hydrogen in form- 
ing salts. In the formation of potassium nitrate from 
nitric acid, HN0 3 , one atom of potassium is substituted 
for the one atom, of hydrogen in the molecule of nitric 
acid, forming the salt KN0 3 . So, also, in sodium nitrate, 
NaN0 3 , one atom of the univalent element sodium is sub- 
stituted for one atom of hydrogen. In the molecule of 
sulphuric acid, H 2 S0 4 , there are two atoms of hydrogen. 
To replace these, two atoms of a univalent element are re- 
quired. Thus, potassium sulphate is K 2 S0 4 , and sodium 
sulphate is Na 2 S0 4 . Illustrations of salts containing 
bivalent metals are the following: Zinc sulphate, ZnS0 4 , 
in which one atom of the bivalent element zinc is sub- 



VARIATIONS IN VALENCE. 205 

stituted for the two atoms of hydrogen in sulphuric acid; 
and barium sulphate, BaS0 4 , in which one atom of bivalent 
barium takes the place of the two atoms of hydrogen in 
sulphuric acid. 

When a bivalent metal forms a salt with an acid like 
nitric acid, which contains but one atom of hydrogen in the 
molecule, it is believed that one atom of the metal acts 
upon two molecules of the acid, thus: 

pn , HN0 3 _ p^ ( N0 3 , TT 

or 

Cu + 2HNO3 = Cu (NO,) a + H 2 . 

The formula of zinc nitrate is similar, viz., Zn(N0 3 ) 2 . 
In the case of trivalent elements the matter is a little more 
complicated, but still simple enough if it is borne in mind 
that a univalent atom replaces one atom of hydrogen; a 
bivalent atom replaces two atoms of hydrogen; a trivalent 
atom replaces three atoms of hydrogen, etc., etc. 

Variations in Valence. — The subject of valence is a diffi- 
cult one to deal with, for the reason that the valence of an 
element is not fixed, but varies according to circumstances. 
It may vary (1) according to the temperature. In general, 
the higher the temperature the lower the valence. Thus, 
phosphorus, which is quinquivalent towards chlorine at 
ordinary temperatures, as is shown by the formation of the 
compound PC1 6 , is trivalent towards the same element at 
higher temperatures, as is shown by the fact that when 
heated the compound PC1 5 gives off chlorine and becomes 
PC1 3 . 

The valence of the element may vary (2) according to 
the chemical character of the element with which it com- 
bines. Thus, phosphorus, which is quinquivalent towards 
chlorine at ordinary temperatures, is trivalent towards 
hydrogen, as is shown by the compound PH 3 . 



206 INTRODUCTION TO CHEMISTRY, 

Generally speaking, however, each element shows a ten- 
dency to act with a particular valence; or if it varies at 
all, the variation is between narrow limits. Nitrogen ap- 
pears as trivalent and quinquivalent; carbon as bivalent 
and quadrivalent, etc. 

Summary. — Avogadro's law that equal volumes of all 
gases under the same conditions of temperature and press- 
ure contain the same number of molecules was suggested 
by a comparison of the weights of equal volumes of gases, 
or their specific gravities, with the combining weights as 
found by analysis. The molecular weights of substances 
bear to one another the same relations as the specific 
gravities of their gases or vapors. Owing to a peculiarity 
of gases and vapors which we cannot discuss here, their 
specific gravities are not exactly proportional to their molec- 
ular weights. They are very nearly so. From the specific 
gravity the molecular weight is calculated, and then by 
analyzing the compounds the molecular weight is deter- 
mined exactly. 

After analyzing the compounds of an element and de- 
termining their molecular weights, the smallest quantity of 
the element that occurs in any of the compounds is taken 
as the atom, and the weight of this quantity as compared 
with the weight of the smallest quantity of hydrogen found 
in any of its compounds taken as unity is the atomic weight 
of the element. 

Elementary gases and vapors are made up of molecules, 
which in turn consist of atoms of the same kind. Ele- 
ments are more active in the nascent state than in the free 
state, probably because the instant they are set free the 
atoms are uncombined, while after they have been set free 
these atoms are combined in the form of molecules. 

Formulas of compounds are. intended to represent the 
composition of molecules and their relative weights, They 



SUMMARY. 207 

rest upon analyses and determinations of the specific gravity 
of the substances in the form of gas or vapor. 

The valence of an element is the property by virtue of 
which its atom has the power to hold in combination a cer- 
tain number of other atoms. Elements are called univa- 
lent, bivalent, trivalent, quadrivalent, etc., according as 
they exhibit the simplest valence like that of hydrogen and 
chlorine, or double, treble, or quadruple this valence. 

The substituting power of the elements is determined by 
the valence. An atom of a univalent element can take 
the place of one atom of hydrogen; an atom of a biva- 
lent element can take the place of two atoms of hydrogen. 



CHAPTER XIV. 

CLASSIFICATION OF THE ELEMENTS.— PERIODIC LAW. 

General. — It is difficult to classify the elements satisfac- 
torily, for the reason that, if one set of properties is made 
the basis of classification, it is questionable whether there 
may not be more fundamental properties which should fur- 
nish the basis. As our knowledge in regard to the funda- 
mental properties of the elements increases, the problem of 
classification will become simpler. 

Acid and Basic Properties. — The chemical properties 
that force themselves upon the attention most promi- 
nently in every field of chemistry are those which are 
known as acid properties and basic properties. As has al- 
ready been pointed out, these two kinds of properties are 
complementary. Whatever developments there may be in 
the study of chemistry in the future, it is certain that the 
distinction between these two kinds of properties will 
always be recognized as important. In general, both acids 
and bases contain oxygen and hydrogen. There are some 
elements whose compounds ivith hydrogen and oxygen have 
basic properties, and others ivhose compounds with hydrogen 
and oxygen have acid properties. This important fact may 
be used as a basis for a partial classification of the elements. 
According to this, we have (1) acid-forming elements and 
(2) base-forming elements. As examples of the first class, 
the elements chlorine, nitrogen, and carbon, already studied, 
may be mentioned. Examples of the second class are so- 
dium, calcium, magnesium^ copper, iron, zinc, etc. The 



NATURAL FAMILIES OF THE ELEMENTS. 209 

last-mentioned elements are generally called metals, and 
the acid-forming elements are generally called non-metals. 
The line between acid-forming and base-forming elements 
cannot be drawn sharply, for there are some elements that 
form both acids and bases, according to the relative quan- 
tity of oxygen with which they are combined. Thus, anti- 
mony forms acids with well-marked properties, and also 
other compounds which neutralize acids, and are therefore 
bases. The same is true of chromium, manganese, and 
some other elements. On the other hand, there are several 
elements that form only acids, and several that form only 
bases; and, further, those which form both acids and bases 
generally show a tendency in one direction. In dealing 
with the elements, then, these differences in properties will 
be taken into account. 

Natural Families of the Elements. — Another important 
fact soon recognized is that the elements fall into families 
according to their general chemical properties, the mem- 
bers of the same family showing striking resemblances 
to one another. Thus, there is the chlorine family, 
which includes, besides chlorine itself, bromine, iodine, and 
fluorine. It will soon be seen that these three elements 
resemble chlorine very closely indeed, so that what has 
already been learned in regard to chlorine will be of great 
assistance in the study of the other members of the family. 
Further, there is the sulphur family, consisting of the 
closely related elements sulphur, selenium, and tellurium; 
the potassium family, consisting of lithium, sodium, potas- 
sium, rubidium, and caesium; the calcium family, consist- 
ing of calcium, barium, strontium; and others. In all 
these cases the resemblance between the members of the 
same family is striking. 

Relations between Atomic Weights of the Elements and 
their Properties. — It has long been known that in many 



210 



1NTBODUCTION TO CHEMISTRY. 



cases there is a connection between the atomic weights of 
the elements and their properties. This is illustrated by 
the natural families of which chlorine, sulphur, sodium, 
and calcium are the best-known members. The members 
of the chlorine family most closely related to it are bro- 
mine and iodine; those of the sulphur family are selenium 
and tellurium. Similarly, sodium and lithium are re- 
lated to potassium; and barium and strontium to calcium. 
The atomic weights of these elements are given in the 
table below : 



Chlorine 


35.4 


Sulphur. . . 


32 


Lithium . . . 


. 7 


Calcium... 


40 


Bromine 


80 


Selenium . . 


79 


Sodium. . . . 


. 23 


Strontium. 


87.5 


Iodine . . . 


127 


Tellurium.. 


125 


Potassium . 


. 39 


Barium . . . 


137 



It will be seen that the atomic weight of bromine is nearly 
the mean of those of the other two members of the family. 
35.4 + 127 



For, 

other families: 



= 81.2. The same relation holds in the 



4/ Z 



and 



40 + 137 



= 88.5. 



Similar relations are met with throughout the list of chem- 
ical elements, and a thorough study of the subject has led 
to the remarkable conclusion that the connection between 
the atomic weights and properties of the elements is gen- 
eral. This was first shown by the Eussian chemist Men- 
deleeff, the German chemist Lothar Meyer, and the Eng- 
lish chemist Newlands. 

The Periodic Law. — If, leaving out hydrogen, and begin- 
ning with lithium, which next to hydrogen has the lowest 
atomic weight, the elements are arranged in the order of 
their atomic weights, the first fourteen exhibit a remarka- 
ble relation, as shown in this table : 

Li =7; Gl = 9; B =11; C =12; N = 14; = 16; F = 19; 
Na = 23; Mg=24; Al = 27; Si = 28; P = 31; S = 32; Cl = 35.4, 



THE PERIODIC LAW. 211 

The elements whose symbols stand in the same vertical 
column in this table have similar chemical properties. The 
resemblance is marked in the case of lithium and sodium; 
carbon and silicon; nitrogen and phosphorus; oxygen and 
sulphur; and fluorine and chlorine. Proceeding in the 
same way, the element with the next higher atomic weight 
is potassium, 39.1. This comes in the same vertical column 
with lithium and sodium or with members of the same 
family. Then follow calcium, scandium, titanium, vana- 
dium, chromium, and manganese, each of which falls nat- 
urally in the vertical column containing elements allied 
to it. It has been found possible in this way to arrange 
all the elements except hydrogen — which, strange to say, 
finds no place in the system — in one table exhibiting the 
relations between their atomic weights and properties. 
Several tables have been proposed, but they do not differ 
essentially from one another. The table on page 212 is a 
simple arrangement. 

When the eighth element in the order of the increasing 
atomic weights is reached it is found that it is very much 
like lithium. It is sodium. If this is placed below lithium, 
and the next six elements are placed in the same horizontal 
line, when the fifteenth element is reached it is found like 
the eighth to resemble lithium. Up to and including man- 
ganese there are twenty-one elements excluding hydrogen. 
These fall naturally into three series of seven members 
each, and arranging the symbols of these horizontally, those 
elements that fall in the same vertical columns have the 
same general character. The three elements following 
manganese, viz., iron, cobalt, and nickel, are very much 
alike, and they certainly do not belong in any one of the 
groups. The next element, copper, has some properties 
which ally it to the members of Group I. The next six 
elements fall in Groups II to VII and are in their proper 
places, and the next six fall in Groups I to VII, and are 
also in their proper places as far as the properties are con- 



212 



INTRODUCTION TO CHEMISTRY. 



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THE PEBIODIC LAW. 213 

cerned. After molybdenum in the sixth series comes a 
blank. There is no element known to fill that place. It 
is, however, probable that there is one undiscovered, with 
the atomic weight approximately 100 and with properties 
similar to those of manganese. Then follow three elements 
which resemble one another as closely as iron, cobalt, and 
nickel. These do not belong in Groups I, II, and III, but 
form a small independent group. These two groups of 
three elements occur at the end of the fourth and sixth 
series respectively. One would therefore naturally expect 
a similar group at the end of the eighth series. No such 
group is known, however, though at the end of the tenth 
series, where one would naturally look for the next similar 
small group, there are the three elements osmium, iridium, 
and platinum. The elements of Series 2, beginning with 
lithium and ending with fluorine, differ in some respects 
quite markedly from the other elements of the groups to 
which they belong, as will be seen later. Beginning with 
sodium, it will be seen that there are two series of seven 
elements and a short series of three; then again two series 
of seven and a series of three ; and, although the following 
series are imperfect, it is easy to recognize that the same 
general arrangement of the elements holds good to the end. 
A series of seven elements is called a short period ; while 
two short periods with the accompanying three similar ele- 
ments constitute a long period. The remarkable relations 
thus presented are summed up in the periodic law : 

The properties of an element are periodic functions of its 
atomic weight. 

Composition of Compounds with Hydrogen and with 
Oxygen.— Passing from left to right in each series, the 
elements combine with a larger and larger relative quan- 
tity of oxygen. The only oxygen compound of lithium 
has the formula Li 2 0. The oxide of glucinum is GIO; 
that of boron, B 2 3 ; the highest oxide of carbon is CO • 



214 INTRODUCTION TO CHEMISTRY. 

that of nitrogen, N 2 5 ; that of sulphur, S0 3 ; and that of 
chlorine, C1 2 7 . On. the other hand, the power to combine 
with hydrogen increases from right to left until a limit is 
reached, as is shown by the formulas FH, 0H 2 , NH 3 , and 
CH 4 . 

Acid-forming and Base-forming Elements. — Those ele- 
ments which have the strongest metallic character, whose 
hydroxides are the strongest bases, are included in Group 
I. The hydroxides of the metals in Group II are weaker 
bases, those of the elements in Group III are weaker still, 
while the hydroxides of some of the elements included in 
Group IV have weak acid properties and no basic proper- 
ties. The elements of Group V are nearly all acid-form- 
ing. Those of Group VI form strong acids, as do those of 
Group VII. 

The Weight of its Atom Determines the Properties of an 
Element. — If the atomic weight of an element is known, 
its position in the table is known, and from its position its 
properties can be stated with considerable accuracy. When 
the table was first constructed, the three elements scandium, 
gallium, and germanium were undiscovered. It was seen, 
however, that the gaps existed, and it was predicted that 
elements would be found with atomic weights approxi- 
mately 44, 69, and 72 respectively, and that these elements 
would have certain properties which were clearly described. 
It was suggested that the element with the atomic weight 
44 would bear to calcium and titanium about the same re- 
lation that aluminium bears to magnesium and silicon. 
The predictions were soon after confirmed, and the de- 
scription of the element given before it was discovered was 
found to be singularly correct. The predictions in regard 
to gallium and germanium were also verified most strik- 
ingly. Unquestionably the properties of the elements are 
determined by their atomic weights. An element whose 
atom weighs 100 times as much as that of hydrogen must 



PLAN TO BE FOLLOWED. 215 

have certain definite properties. It must combine with 
hydrogen and with oxygen in certain proportions; it must 
be allied to the members of the chlorine family; its prop- 
erties are the result of that particular weight. Farther, 
it seems to follow that the elements are not entirely in- 
dependent forms of matter, but that they are in all proba- 
bility compounds of a small number of simple elements at 
present unknown to us. Of this, however, there is no 
evidence, and until some one succeeds in isolating one or 
more of these subtler elements it is almost useless to specu- 
late in regard to them. 

Plan to be Followed. — The elements hydrogen, oxygen, 
chlorine, nitrogen, and carbon have been studied in order 
to illustrate the methods of studying chemical problems in 
general, and as examples of the chemical elements. Now, 
following the suggestions of the periodic law, a number of 
other elements will be treated as members of families or 
groups. Hydrogen does not belong to any group. Oxy- 
gen has peculiarities that distinguish it from most other 
elements, but it nevertheless resembles sulphur in many 
ways, and the two are treated together. Chlorine, as 
already stated, belongs to a group of which fluorine, 
bromine, and iodine are the other members. Nitrogen 
belongs to a group of which phosphorus, arsenic, and an- 
timony are the other best-known members. Carbon also 
belongs to a group, silicon being the other well-known 
member. We therefore have the following groups first to 
deal with : 

Chlorine Group. Sulphur Group. Nitrogen Group. Carbon Group. 



Chlorine, 


Sulphur, 


Nitrogen, 


Carbon, 


Bromine, 


Selenium, 


Phosphorus, 


Silicon. 


Iodine, 


Tellurium. 


Arsenic, 




Fluorine. 




Antimony, 
Bismuth. 





The principal members of these groups are acid-forming 
elements. They are generally called non-metals. In the 



216 INTRODUCTION TO CHEMISTRY. 

nitrogen group, however, two of the members are both acid- 
forming and base-forming. There is a gradation in the 
properties from nitrogen to bismuth. 

As the object of this book is to present concisely such 
facts as serve to illustrate the general character of chemical 
action and the general principles of the science of chemis- 
try, it will not be necessary to go into details in dealing 
with these groups. One member of each group having 
been treated comparatively fully, the other members may 
be treated briefly. It w T ill thus be possible to get a clearer 
idea of the principles of the science than by attempting to 
study a large number of facts the connection betweei: 
which can be but dimly discerned, if discerned at all. 

After the acid-forming elements have been studied, the 
base-forming elements will be taken up in a similar way; 
but, as will be seen, the chemistry of the acid-forming ele- 
ments exhibits more variety, and is hence better adapted 
to the illustration of the general principles of the science 
than that of the base-forming elements, so that the latter 
need not be treated as fully. 



CHAPTER XV. 

THE CHLORINE GROUP : 
CHLORINE, BROMINE, IODINE, FLUORINE. 

The three members of this group which show the most 
marked resemblance are chlorine, bromine, and iodine. 
Fluorine has properties of the same general kind, and its 
compounds resemble those of the other three members of 
the group, so that it is properly treated with them. 

Bromine, Br (At. Wt. 80). — This element occurs in 
nature in company with chlorine. Chlorine, as has been 
stated, occurs mostly in combination with sodium, as 
sodium chloride, or common salt. In several of the great 
salt-beds there is some bromine in the form of sodium 
bromide, NaBr, and in some places it occurs as potassium 
bromide, KBr. 

Preparation. — The process of preparation of bromine is 
exactly the same as that made use of for extracting chlo- 
rine. It will be remembered that in order to get chlorine 
from sodium chloride the salt is first converted into hydro- 
chloric acid, and this is then oxidized. So, too, in order 
to get bromine from sodium bromide, it must first be con- 
verted into hydrobromic acid, and this then oxidized. The 
reactions involved are usually : 

SNaBr + H 2 S0 4 = NaJS0 4 + 2HBr; 
2HBr + = H 2 + 2~Br. 

217 



218 INTRODUCTION TO CHEMISTRY. 

As in the case of chlorine, the substance commonly used 
is manganese dioxide, when the reaction takes place accord- 
ing to the following equation : 

4HBr + Mn0 2 = MnBr 2 + 2H 2 + 2Br. 

[Eefer back to the explanation of this reaction given 
under the head of chlorine. What other methods might 
be used in the preparation of bromine ?] 

Properties. — Bromine is a heavy, dark red liquid at ordi- 
nary temperatures. It is easily converted into vapor which 
is brownish red. At — 7.3° it is solid. It has an ex- 
tremely disagreeable smell, to which fact it owes its name 
(from fipGo/tos, a stench). 

Its properties are, in general, like those of chlorine. It 
acts violently upon organic substances. It attacks the 
skin and the membranes lining the passages of the throat 
and lungs in much the same way as chlorine. Wounds 
caused by the liquid coming in contact with the skin are 
painful and serious. It must be handled with great care. 
With water at low temperatures it forms a hydrate cor- 
responding to chlorine hydrate, of the formula Br 2 .10H 2 O, 
which decomposes when left in contact with the air at 
ordinary temperatures. It dissolves slightly in water, form- 
ing a colored solution called bromine-water. 

Its chemical conduct is also like that of chlorine. It 
combines with many elements directly and with great avid- 
ity. Its combination with arsenic and some other elements 
is accompanied by an evolution of light and heat, as in the 
case of chlorine. Its compounds with other elements are 
called bromides. While acting in general in the same way 
as chlorine, it is a somewhat weaker element, so that chlo- 
rine drives it out of its compounds and sets it free. 

Experiment 103. — Mix together 3.5 grams potassium bromide 
and 7 grams manganese dioxide. Put the mixture into a 500-cc. 



HYDROBROMIC ACID. 219 

flask ; connect with a condenser (see Fig. 25). Mix 15 cc. con- 
centrated sulphuric acid and 90 cc. water. After the liquid is 
cool pour it upon the mixture in the flask. Heat gently, when 
bromine will be given off in the form of vapor. A part of this will 
condense and collect in the receiver. Perform this experiment 
under a hood with a good draught. In treating the manganese 
dioxide and potassium bromide together with sulphuric acid, the 
action takes place as represented in the following equation : 

2KBr + MnO, + 2H 2 S0 4 = K 2 S0 4 + MnSO* + 2H 2 + Br* 

Hence both potassium sulphate, K 2 S0 4 , and manganese sulphate, 
MnS0 4 , are left behind in the flask. 

[When sulphuric acid acts upon manganese dioxide the action 
takes place thus : 

Mn0 2 + H 2 S0 4 == MnS0 4 + H 2 + O. 

If this action took place in the presence of hydrobromic acid, 
what effect would the liberated oxygen have ? Suppose the oxygen 
were allowed to escape from the flask containing the manganese 
dioxide and sulphuric acid, and then passed into hydrobromic 
acid, would the same result be reached as when the hydro- 
bromic acid is in the flask from which the oxygen is liberated ? 
What is the commonly accepted explanation ? If the formula of 
manganese sulphate is MnS0 4 , what is the valence of manganese ? 
What would you expect the formula of manganese chloride to be ? 
Of manganese oxide ? Is the valence of manganese greater 
toward oxygen or toward chlorine ?] 

Hydrobromic Acid, HBr. — The only compound that 
bromine forms with hydrogen alone is hydrobromic acid. 
This is in all respects very much like hydrochloric acid. It 
is made in the same w T ay. It is a colorless gas which forms 
fumes in the air in consequence of its attraction for moist- 
ure. Its solution in water acts very much like ordinary 
hydrochloric acid. The elements are not held together as 
firmly in hydrobromic as in hydrochloric acid. This is 
shown by its decomposition under circumstances in which 
hydrochloric acid is stable. Thus, for example, it is de- 



220 INTRODUCTION TO CHEMISTRY. 

composed by sulphuric acid, while hydrochloric acid is not. 
The hydrogen is separated from the bromine and acts upon 
the sulphuric acid, while the bromine is given off as such. 
Hence, when potassium bromide is treated with sulphuric 
acid, hydrobromic acid is given off, together with bromine 
and a compound of sulphur and oxygen which is formed 
by the action of hydrogen on the sulphuric acid. 

Experiment 104. — In a small porcelain evaporating-dish put a 
few crystals of potassium bromide. Pour upon them a few drops 
of concentrated sulphuric acid. The white fumes of hydrobromic 
acid and the reddish-brown vapor of bromine are noticed. Treat 
a few crystals of potassium or sodium chloride in the same way. 
What difference is there between the two cases ? 

Compounds with Hydrogen and Oxygen. — With hydrogen 
and oxygen bromine forms compounds that resemble very 
closely those which chlorine forms with the same elements. 
The principal ones are iromic and liypohromous acids. 
The potassium salt of bromic acid, j^Br0 3 , is formed by 
treating a strong solution of caustic potash with bromine : 

3Br 2 + 6KOH = 5KBr + KBr0 3 + 3H 2 0. 

The potassium salt of * hypobromous acid, J^ferO, is 
formed by treating a dilute solution of caustic potash with 
bromine : 

Br 2 + 2KOH = KBr + KBrO + H 2 0. 

Iodine, I (At. Wt. 127). — This element occurs in nature 
in combination with sodium, in company with chlorine and 
bromine, but in smaller quantity than either. It is also 
found in larger quantities in all sea-plants. It is obtained 
largely from the latter source. On the coasts of Scotland 
and France the sea- weed which is thrown up by storms is 
gathered, dried, and burned. The organic portions are thus 
destroyed [What is the meaning of the word destroyed used 
in this sense ?] and the mineral or earthy portions are left 



IODINE. 221 

behind as ashes. This incombustible residue, which is called 
kelp, contains sodium iodide. Sea-weed is also cultivated for 
the sake of the sodium iodide contained in it. Chili salt- 
petre, or the sodium nitrate found in Chili, contains some 
sodium iodide, and this now furnishes a considerable quan- 
tity of the iodine of commerce. 

Iodine is obtained from sodium iodide just as chlorine 
and bromine are obtained from their compounds with 
sodium and potassium. [Give the equations representing 
the steps which must be taken in order to separate iodine 
from sodium iodide.] 

Experiment 105. — Mix together about 2 grams of sodium or 
potassium iodide and 4 grams manganese dioxide. Treat with a 
little sulphuric acid in a one to two litre flask. Heat gently on a 
sand-bath. Gradually the vessel will be filled with the beautiful 
colored vapor of iodine. In the upper parts of the flask some of 
the iodine will be deposited in the form of crystals of a grayish- 
black color. 

Properties. — At ordinary temperatures iodine is a gray- 
ish-black crystallized solid. It is volatile at ordinary tem- 
peratures. It acts upon the mucous membranes, though 
less energetically than chlorine and bromine. It colors the 
skin yellowish brown, and acts as an absorbent, causing the 
reduction of swellings. It melts at 113-115°, and boils at 
250°, when it is converted into a violet vapor. 

The action of iodine is, in general, the same as that of 
chlorine and bromine, only its affinities are weaker. Hydro- 
bromic acid, as we have seen, is a weaker compound than 
hydrochloric acid. Hydriodic acid is still weaker. Chlorine 
acting upon hydrobromic acid sets bromine free. Chlorine 
and bromine set iodine free from hydriodic acid. 

Iodine dissolves slightly in water, easily in alcohol, and 
easily in a water-solution of potassium iodide. 

Experiment 106. — Make solutions of iodine in water, in alco- 
hol, and in a water-solution of potassium iodide. Use small 
quantities in test-tubes, 



222 INTRODUCTION TO CHEMISTRY. 

When a solution containing free iodine is treated with a 
little starch-paste, the solution turns blue, in consequence 
of the formation of a complicated compound of starch and 
iodine. Bromine and chlorine do not form blue com- 
pounds. Advantage is taken of this fact to distinguish 
between iodine and other members of the same family. 

Experiment 107. — Make some starch-paste by covering a few 
grains of starch in a porcelain evaporating-dish with cold water, 
grinding this to a paste, and pouring 200-300 cc. boiling-hot 
water on it. After cooling add a little of this paste to a dilute 
water-solution of iodine. The solution will turn blue if the con- 
ditions are right. Now add a little of the paste to a dilute water- 
solution of potassium iodide. There is no change of color, because 
the iodine is in combination with the potassium. Add a drop or 
two of a solution of chlorine in water, when the blue color will 
appear. The explanation of this phenomenon is that the chlorine 
sets the iodine free, and the free iodine then acts upon the starch, 
producing the blue compound. [How can you show that the 
chlorine itself will not form a blue compound with starch ?] 

Hydriodic acid, HI, is analogous to hydrochloric and 
hydrobromic acids. It is set free from its compounds by 
treating them with sulphuric acid, but it is even more un- 
stable than hydrobromic acid, and hence breaks up into 
hydrogen and iodine. The iodine is liberated, while the 
hydrogen acts on the sulphuric acid, as it does in the case 
of hydrobromic acid. 

Experiment 108. — Treat a few crystals of potassium iodide 
with sulphuric acid. [What do you notice ?] Compare the result 
with that obtained in the case of potassium bromide and sodium 
chloride. 

Iodic Acid, HI0 3 . — The principal compound of iodine 
with hydrogen and oxygen is iodic acid, HI0 3 , which cor- 
responds to chloric and bromic acids. It is known princi- 
pally in the form of its potassium salt, potassium iodate, 
EJ0 3 . When heated, this salt, like the chlorate and the 



FLUORINE— HYDROFLUORIC ACID. 223 

bromate, gives up all its oxygen, potassium iodide, KI, 
being left behind. 

Fluorine, F (At. Wt. 19). — This element occurs in 
nature in large quantity, and widely distributed, but al- 
ways in combination with other elements. It is found 
chiefly in combination with calcium, as fluor-spar, or 
calcium fluoride, CaF 2 , and in combination with sodium 
and aluminium, as cryolite, a mineral which occurs abun- 
dantly in Greenland and has the composition 3NaF.AlF 3 , 
being a complex compound of sodium fluoride and alumin- 
ium fluoride. 

All attempts to obtain fluorine in the free state failed 
until a few years ago, when its isolation was effected by 
passing an electric current through liquid hydrofluoric 
acid in a platinum vessel. 

Properties. — Fluorine is the most active of all the ele- 
ments at ordinary temperatures. It is a colorless gas. It 
acts upon almost all substances. Thus, it decomposes 
water, yielding ozone and hydrofluoric acid; at ordinary 
temperatures it combines directly with sulphur, phosphorus, 
iron, etc., with evolution of light and heat. It does not, 
however, act upon platinum. 

Hydrofluoric acid, HF, is made from fluor-spar by treat- 
ing it with sulphuric acid. The action is of the same kind 
as that which takes place when hydrochloric acid is liber- 
ated from sodium chloride : 

CaF 2 + H a S0 4 = CaSO, + 2HF. 

It is a colorless gas, with strong acid properties. It 
greatly irritates the membranes lining the respiratory or- 
gans, and hence care should be taken not to inhale it. It 
acts upon glass, dissolving it, and must therefore be kept 
in vessels of rubber, lead, or platinum, upon which it does 
not act. Its action on glass consists in the transformation 



224 INTRODUCTION TO CHEMISTRY. 

of silicon dioxide, or silica, Si0 2 , which is contained in all 
kinds of glass, into silicon tetrafluoride, SiF 4 , which is a 
gas. The action is represented thus : 

Si0 2 + 4HP = SiF 4 + 2H 2 0. 

Experiment 109. — In a lead or platinum vessel put a few grams 
(5-6) of powdered fluor-spar and pour upon it enough concentrated 
sulphuric acid to make a thick paste. Cover the surface of a 
piece of glass with a thin layer of wax or paraffin, and through 
this scratch some letters or figures, so as to leave the glass exposed 
where the scratches are made. Put the glass over the vessel con- 
taining the fluor-spar, and let it stand for some hours. Take off 
the glass, scrape off the coating, and the figures which were 
marked through the wax or paraffin will be found etched on the 
glass. 

The acid is used for etching glass, particularly for mark- 
ing scales on thermometers, barometers, and other gradu- 
ated glass instruments. A solution of the gas in water is 
manufactured for this purpose and kept in rubber bottles. 

Fluorine does not combine with oxygen. It is the only 
element of which this statement is true. 

Comparison of the Members of the Chlorine Group. — In 

considering, first, the physical properties of these elements, 
we notice that all, with the exception of fluorine, form 
colored gases or vapors. At ordinary temperatures chlo- 
rine is a gas, bromine a liquid, and iodine a solid. In re- 
gard to their chemical conduct, it may be said that, in gen- 
eral, fluorine is the most energetic; chlorine comes next in 
order, then bromine, and lastly iodine. This is seen par- 
ticularly in the relative stability of their compounds with 
hydrogen. Their compounds with metals also show the 
same relation. On the other hand, with oxygen the order 
is reversed. Fluorine does not unite with oxygen at all. 
The compounds of chlorine and oxygen are very unstable; 
those with bromine rather more stable; and one compound 
of iodine and oxygen is comparatively stable. 






MEMBERS OF THE CHLORINE GROUP COMPARED. 225 

The elements of this group combine with hydrogen and 
with other elements in the simplest way. They are all 
univalent. 

The compounds formed by the three elements chlorine, 
bromine, and iodine with hydrogen and oxygen have analo- 
gous composition, and are formed by analogous reactions. 
Thus, we have the hydrogen compounds : 

HC1, HBr, and HI; 
and the compounds with hydrogen and oxygen: 

HCIO HBrO 

HCIO. 



HCIO, HBr0 3 HIO3 
HC10 4 HI0 4 

The properties of any compound of one element are 
similar to those of the compounds of analogous composi- 
tion of the other elements of the group. 

All these facts seem to indicate that these elements are 
not distinct forms of matter entirely independent of one 
another, but rather that they contain some common con- 
stituent. The relations between the atomic weights of the 
members of the group have already been referred to. 



CHAPTER XVI. 

THE SULPHUR GROUP: 
SULPHUR, SELENIUM, TELLURIUM. 

Sulphur, S (At. Wt. 32). — The principal member of this 
group is sulphur. In nature it is frequently found accom- 
panied by small quantities of selenium, and sometimes by 
tellurium. It has been known in the elementary form 
from the earliest times, for the reason that it occurs abun- 
dantly in this form in nature. It is found particularly in 
the neighborhood of volcanoes, as in Sicily, which is the 
chief source of the sulphur of commerce. It occurs, fur- 
ther, in combination with many metals as sulphides, — as in 
iron pyrites, FeS 2 ; copper pyrites, FeCuS 2 ; galenite, PbS, 
etc.; in combination with metals and oxygen as sulphates, 
— for example, as calcium sulphate, or gypsum, CaS0 4 -f- 
2H 2 0; barium sulphate, or heavy spar, BaS0 4 ; lead sul- 
phate, PbS0 4 ; in a few vegetable and animal products in 
combination with carbon, hydrogen, and, generally, with ni- 
trogen. 

Extraction of Sulphur from its Ores. — When taken from 
the mines, sulphur is mixed with many earthy substances 
from which it must be separated. This separation is ac- 
complished by piling the ore in such a way as to leave pas- 
sages for air. The piles are covered with material to pre- 
vent free access of air, and the mass is then lighted below. 
A part of the sulphur burns, and the heat thus furnished 
melts the rest of the sulphur. The molten sulphur runs 
down to the bottom of the pile, and is drawn off from time 

22a 



SULPHUR. 227 

to time. If the pile were not protected from free access of 
air, the sulphur would burn up, yielding a gas, sulphur 
dioxide, S0 2 . 

[What analogy is there between this process and that 
employed in making charcoal ? What are the essential 
differences between the two processes ?] 

Refining of Sulphur. — The crude brimstone thus obtained 
is afterwards refined by distillation, and it is this distilled 
sulphur which comes to market under the names "roll 
brimstone" and "flowers of sulphur." The distillation is 
carried on in earthenware retorts connected with large 
chambers of brick-work. When the vapor of sulphur first 
comes into the condensing-chamber it is suddenly cooled, 
and hence deposited in the form of a fine powder. This is 
called "flowers of sulphur." After the distillation has 
continued for some time, the vapor condenses in the form 
of a liquid, which collects at the bottom of the chamber. 
This is drawn off into wooden moulds and takes the form 
of "roll brimstone" or "stick sulphur." 

Properties. — Sulphur is a yellow, brittle substance which 
at — 50° is almost colorless. It melts at 114.5°, forming a 
thin, straw-colored liquid. When heated to a higher tem- 
perature it becomes darker and darker in color, and at 200° 
to 250° it is so viscid that the vessel in which it is con- 
tained may be turned upside down without danger of its 
running out. Finally, at 448.4° it boils and is then con- 
verted into a brownish-yellow vapor. 

Experiment 110. — Distil about 10 grams of roll sulphur from 
an ordinary glass retort. Notice the changes above described; 
Collect the liquid sulphur which passes over, in a beaker-glass 
containing cold water. » 

Crystals of Sulphur. — When molten sulphur solidifies, or 
when it is deposited from a solution, its particles arrange 



228 INTRODUCTION TO CHEMISTRY. 

themselves in regular forms called crystals. But, strange 
to say, the crystals formed from molten sulphur are entirely 
different from those deposited from solutions of sulphur. 
The former are honey-yellow needles. The latter are octa- 
hedrons with rhombic base, which is also the form of the 
sulphur found in nature. A careful examination of the 
needles shows that the angles which their faces form with 
one another are not the same as the angles formed by the 
faces of the octahedrons, and that the crystals are con- 
structed on a different plan. The needles belong to the 
monoclinic system of crystals, and the octahedrons to the 
rhombic system. 

Crystallography. — Notwithstanding the infinite number 
of forms assumed by solids in passing from the liquid to the 
solid state and when deposited from solutions, it has been 
shown that all can be referred to a very few systems. 
Usually six systems are adopted. These are : 

1. The Regular System. All the crystals belonging to 
this system can be referred to three axes of equal length, 
and at right angles to one another, crossing at the centre. 
Examples of crystals belonging to this system are the regu- 
lar octahedron and the cube. The three axes are the 
imaginary lines which pass through the solid angles of the 
octahedron. All the other forms of this system may be 
referred to this octahedron. 

2. The Tetragonal System. In this the forms are 
referred to three axes at right angles, two of equal length 
and one differing from the other two. The fundamental 
forms are the octahedron and prism. 

3. The Hexagonal System. The crystals of this sys- 
tem are referred to four axes, — three of equal length in- 
clined at 60° to one another, and a fourth at a right angle 
to them, which is either of the same length or of a different 
length. The six-sided pyramid and prism are the principal 
forms, 



CRYSTALLOGRAPHY. 229 

4. The Khombic System. The crystals belonging to 
this system have three axes, of unequal lengths, at right 
angles to one another. 

5. The Monoclls'ic System. In this system the crys- 
tals have three axes, — two at right angles to each other, 
the third at right angles to one and inclined to the other. 

6. The Triclixic System. The crystals belonging to 
this system are referred to three axes, all inclined to one 
another. 

The subject of crystallography is one that cannot be 
made clear in a few words. It requires careful study and 
much practice in observing forms of crystals. From what 
has just been said, however, it will be seen that the system 
of classification of crystals is a simple one. For our pres- 
ent purpose, the fact should be specially emphasized that 
the crystalline form of a substance is a very definite prop- 
erty, by means of which it may be distinguished from other 
substances. The fact that a substance crystallizes in the 
regular system is just as characteristic of that substance as 
the fact that it boils or melts at a certain point. Thus, we 
know that ice always melts at 0°, and that water solidifies 
at 0°. We should be much surprised to find water solidi- 
fying at some ^other temperature, say 20°. Similarly, 
knowing that sulphur occurs in nature crystallized in 
forms which belong to the rhombic system, we are natu- 
rally surprised to find that, when molten sulphur solidifies, 
it crystallizes in forms belonging to the monoclinic system. 
What is perhaps still stranger is the fact that when the 
honey-yellow needles are allowed to stand unmolested they 
spontaneously undergo a change. They become opaque; 
their color changes; and now, if examined carefully, they 
are found to consist of minute crystals like those found in 
nature. It is evident that the arrangement of the particles 
in the monoclinic crystals of sulphur is not a stable one. 

Substances which crystallize in two distinct forms are 



230 INTRODUCTION TO CHEMISTRY. 

called dimorphous. Carbon crystallizes in two different 
forms [What are they ?], and is hence dimorphous. 

Experiment 111. — In a covered porcelain crucible melt a few 
grams of roll sulphur. Let it cool slowly, and when a thin crust 
has formed on the surface make a hole through this and pour out 
the liquid part of the sulphur. The inside of the crucible will be 
found lined with the honey-yellow needles which, as has been 
stated, belqng to the monoclinic system. Take out a few of the 
crystals and examine them. Are they brittle or elastic ? What 
is their color ? Are they opaque, transparent, or translucent ? 
Lay the crucible aside, and in the course of a few days again 
examine the crystals. What changes, if any, have taken place ? 

Other Forms of Sulphur. — Sulphur can also be obtained 
in the amorphous, or uncrystallized, condition. That which 
was collected under water in Experiment 110 will be found 
to be soft and doughlike. It is amorphous. After a time 
it becomes brittle. When separated from a compound 
which is dissolved in water, it is finely divided, and gives 
the liquid an appearance suggesting milk. 

Crystallization from Carbon Bisulphide. — Sulphur is in- 
soluble in water, slightly soluble in alcohol and ether. It 
dissolves in carbon bisulphide, CS 2 , and from the solution 
it is deposited in rhombic crystals. 

Experiment 112. — Dissolve 2 to 3 grams roll sulphur in 5 to 
10 cc. carbon bisulphide. Put the solution in a shallow vessel, 
and allow the carbon bisulphide to evaporate by standing in the 
air. The sulphur will remain behind in the form of crystals. 

Chemical Conduct of Sulphur. — Sulphur combines with 
oxygen when heated to a sufficiently high temperature, 
forming sulphur dioxide, S0 2 . [Compare carbon and sul- 
phur in this respect.] It combines readily with most 
metals, forming sulphides, which are analogous to the 
oxides. Its combination with iron has already been shown 
in Experiment 10. It also combines with copper, the act 
being accompanied by light and heat. 



HYDROGEN SULPHIDE. 231 

Experiment 113. — In a wide test-tube heat some sulphur to 
boiling. Introduce into it small pieces of copper-foil or sheet- 
copper. Or hold a narrow piece of sheet-copper so that the end 
just dips into the boiling sulphur. 

Hydrogen Sulphide, Sulphuretted Hydrogen, H 2 S. — When 
hydrogen is passed over highly-heated sulphur, the two 
elements combine to form hydrogen sulphide. [Is there 
any analogy between this process and the formation of 
water by the burning of hydrogen?] This compound of 
sulphur and hydrogen occurs in nature in solution in the 
so-called sulphur waters, whicn are met with in many parts 
of this and other countries. It also issues from the earth 
in some places. It is formed by heating organic substances 
which contain sulphur, just as water is formed by heating 
organic substances which contain oxygen, and ammonia by 
heating such as contain nitrogen. It is formed, further, 
by decomposition of organic substances which contain 
sulphur, as, for example, the albumin of eggs. The odor 
of rotten eggs is partly due to the formation of hydrogen 
sulphide. 

Preparation. — In the laboratory the gas is most readily 
made by treating a sulphide with an acid. When a metal, 
as iron, is treated with sulphuric acid, hydrogen is given 
off and the iron salt of the acid is formed thus: 

Fe + H 2 S0 4 = FeS0 4 + H 2 . 

When sulphuric acid acts upon the' oxide of iron, hydro- 
.a is given off in combination with oxygen as water, thus: 

FeO -f H 2 S0 4 = FeS0 4 + H 2 0. 

Finally, when sulphuric acid acts upon iron sulphide, 
hydrogen is given off in combination with sulphur as hy- 
drogen sulphide, thus: 

FeS + H 2 S0 4 = FeS0 4 + H 2 S. 



INTRODUCTION TO CHEMISTRY. 



A similar explanation holds for other acids. For exam- 
ple, hydrochloric acid acts upon iron sulphide in accord- 
ance with the equation 

2H01 + FeS = FeCl 2 + H 2 S. 

Experiment 114. — Arrange an apparatus as shown in Fig. 51. 
Put a small handful of iron sulphide, FeS, in the flask, and 
pour dilute sulphuric acid upon it. Pass the evolved gas through 
a little water contained in the wash-cylinder A. Pass some of 




Fig. 51. 

the gas into water. [What evidence have you that it dissolves ?] 
Collect some by displacement of air. Its specific gravity is 1.178. 
[Should the vessel be placed with the mouth down or up ?] Set 
fire to some of the gas contained in a cylinder. If there is fiee 
access of air, the sulphur burns to sulphur dioxide and the 
hydrogen to water. 

Properties. — Hydrogen sulphide is a colorless, transpar- 
ent gas of specific gravity 1.178. It has an extremely dis- 
agreeable odor, somewhat suggestive of that of rotten eggs. 
It is poisonous, even small quantities causing headache, 
vertigo, nausea, and other bad symptoms. It is soluble in 
water, about three volumes being taken up at ordinary 
temperatures. This solution is used in the laboratory 



HYDROGEN SULPHIDE. 233 

instead of the gas. Hydrogen sulphide is easily decom- 
posed into its elements. In consequence of this instability, 
it causes a number of changes which the analogous com- 
pound water cannot effect. The relations here are similar 
to those which exist between hydrochloric and hydriodic 
acids. Hydrochloric acid is very stable, while hydriodic 
acid breaks dow T n readily into hydrogen and iodine. Chlo- 
rine, bromine, and iodine act upon hydrogen sulphide, 
setting the sulphur free and combining with the hydrogen. 
Thus, with chlorine the action takes place as represented 
in the equation 

H 2 S + Cl 2 = 2HC1 + S. 

[Does chlorine ever act in a similar way on water ? Un- 
der what circumstances ? What is the peculiarity of the 
oxygen given off ?] 

Most metals when heated in the gas are converted into 
sulphides. Thus, when it is passed over heated iron this 
reaction takes place : 

Fe + H 2 S = FeS + H 2 . 

[What takes place when water- vapor is passed over 
heated iron ?] 

Many of the sulphides are insoluble in water. Hence, 
when hydrogen sulphide is passed through solutions con- 
taining metals in the form of soluble salts, the insoluble 
sulphides are thrown down, or precipitated. 

Experiment 115.— Pass hydrogen sulphide successively through 
solutions containing a little lead nitrate, zinc sulphate, and 
arsenic prepared by dissolving a little white arsenic, or arsenic 
trioxide, As 2 3 , in dilute hydrochloric acid. In the vessel con- 
taining the lead a black precipitate of lead sulphide will be 
formed ; in the one containing the zinc sulphate there will be 
formed a white precipitate of zinc sulphide ; in the one containing 
the arsenic, a straw-yellow precipitate of arsenic sulphide will be 
formed. In all these cases the hydrogen of the hydrogen sul- 



234 INTRODUCTION TO CHEMISTRY. 

phide and the metal of the salt exchange places. For example, 
in the case of zinc sulphate the reaction takes place thus : 

ZnS0 4 + H 2 S = ZnS + H 2 S0 4 . 

Chemical Analysis. — In dealing with chemical substances 
the first thing we have to determine is their composition, 
or, in other words, we have to analyze them. For this 
purpose the properties of the elements and their general 
conduct towards chemical substances must, of course, be 
known. To facilitate the process of analysis the mixture 
to be examined is usually brought into solution and then 
treated successively with certain substances, the effect 
being observed in each case. Suppose we had a solution 
containing most of the metallic elements in the form of 
salts. If we were to pass through this solution hydrogen 
sulphide, some of the metals would be precipitated in the 
form of sulphides, while others would remain in solution, 
as their sulphides are soluble. We then filter off the pre- 
cipitate and examine it by other methods, and we could 
also further examine the solution from which the sulphides 
were precipitated. By adding to this another reagent 
which will precipitate some of the metals and leave the 
others in solution, we learn still more in regard to the 
composition of the substance under examination. Hydro- 
gen sulphide is constantly made use of in the laboratory 
for the purposes of analysis. 

Hydrosulphides. — When hydrogen sulphide acts upon 
hydroxides, the action consists in the formation of hydro- 
sulphides. In the case of potassium hydroxide the action 
takes place thus : 

KOH + H 2 S = KSH + H,0. 

The oxygen and sulphur simply exchange places. 

If only half enough hydrogen sulphide is passed into the 



SULPHUR DIOXIDE. 235 

solution to effect the above change, a sulphide is formed 
thus : 

2KOH + H 2 S = K 2 S + 2H a O. 

Or if hydrogen sulphide is allowed to act on potassium 
sulphide, the product is potassium hydrosulphide: 

K 2 S + H 2 S = 2KSH. 

Compounds of Sulphur with Oxygen and with Hydrogen 
and Oxygen. — When sulphur burns in the air it forms the 
dioxide S0 2 . Under certain conditions the dioxide com- 
bines with more oxygen, forming the trioxide S0 3 . When 
sulphur dioxide acts upon water, sulphurous acid is formed : 

50 2 + H 2 = H 2 S0 3 . 

[What analogy is there between the acid thus formed 
and carbonic acid ?] 

When the trioxide combines with water, sulphuric acid 
is formed : 

50 3 + H 2 = H 2 S0 4 . 

Sulphur Dioxide, S0 2 . — This compound is formed by 
burning sulphur in the air or in oxygen. It issues from 
volcanoes in large quantities. It is best prepared by treat- 
ing copper with sulphuric acid. The action does not take 
place without the aid of heat. The copper appears first 
to reduce the acid, forming sulphurous acid, H 2 S0 3 , and 
copper oxide, CuO, the former compound breaking down at 
once, however, into sulphur dioxide and water, and the 
latter dissolving in sulphuric acid as sulphate. These 
changes are represented by the three equations: 

(1) H 2 S0 4 + Cu = H 2 S0 3 + CuO ; 

(2) H 2 S0 3 = H 2 + S0 2 ; 

(3) H 2 S0 4 + CuO = CuS0 4 + H 2 0. 

Or combined in one equation these may be represented thus : 
2H 2 S0 4 + Cu = CuS0 4 + 2H 2 + S0 2 . 



236 INTRODUCTION TO CHEMISTRY. 

[Compare the action of copper on sulphuric acid with 
that of copper on nitric acid. What analogy is there be- 
tween the two cases ? What difference ?] 

Sulphur dioxide is a colorless gas of an unpleasant, suffo- 
cating odor, familiar to every one as that of burning sul- 
phur-matches. Water absorbs it readily. 

Experiment 116. — Put eight or ten pieces of sheet-copper, one 
to two inches long and about half an inch wide, in a 500-ce. 
flask; pour 15 to 20 cc. concentrated sulphuric acid upon it. On 
heating, sulphur dioxide will be evolved. The moment the gas 
begins to come off, lower the flame, and keep it at such a height 
that the evolution is regular and not too active. Pass some of the 
gas into a bottle containing water. Fill a vessel by displacement 
of air. Its specific gravity is 2.24. See whether the gas will 
burn or support combustion. 

Sulphurous Acid, H 2 S0 3 . — The solution in water has acid 
properties, and probably contains the acid H 2 S0 3 . By 
neutralizing the solution with bases, the sulphites, or salts 
of sulphurous acid, are obtained. The sulphites are anal- 
ogous to the carbonates in composition, and suffer the same 
decomposition when treated with acids. When a carbonate 
is treated with an acid, carbon dioxide is given off. So, 
also, w T hen a sulphite is treated with an acid, sulphur diox- 
ide is given off : 

Na 2 S0 3 + H 2 S0 4 = Na.,S0 4 + H 2 + S0 2 , 
Na 2 S0 3 + 2HC1 = 2NaCl + H 2 + S0 2 . 

When a solution of sulphur dioxide is allowed to stand in 
the air in loosely-stoppered bottles, it takes up oxygen, the 
sulphurous acid being converted into sulphuric acid: 

H 2 S0 3 + = H 2 S0 4 . 

Sulphur dioxide is a good bleaching agent, and is exten- 
sively used for the purpose of bleaching wool, silk, straw, 
paper, etc. In some cases the bleaching is due to the fact 



SULPHURIC ACID. 237 

that the sulphur dioxide extracts oxygen from the colored 
substances, forming colorless products. In other cases the 
action is more complicated. 

Sulphur dioxide has the power to check fermentation, 
and is used to preserve liquids that have a tendency to 
undergo fermentation. 

Its principal use is in the manufacture of sulphuric acid. 
For this purpose it is made in enormous quantities. 

Experiment 117. — Burn a little sulphur iu a porcelain crucible 
under a bell-jar. Place over the crucible on a tripod some flow- 
ers. In the atmosphere of sulphur dioxide the flowers will be 
bleached. 

Sulphuric Acid, H 2 S0 4 . — Sulphuric acid is found in nature 
in the form of salts, as gypsum, heavy spar, etc. It cannot 
easily be prepared from its salts, as nitric acid and hydro- 
chloric acids are, and is made exclusively by oxidizing sul- 
phur dioxide in the presence of water, or, in other words, 
by oxidizing sulphurous acid. The reactions involved in 
the manufacture of sulphuric acid are: 

s + o 2 =S0 8 , 

S0 2 + H 2 = H 2 S0 3 , 

H 2 S0 3 + = H 2 S0 4 . 

Manufacture of Sulphuric Acid. — The last reaction cannot 
readily be effected directly by the action of the oxygen of 
the air, but an extremely interesting method has been de- 
vised by which the oxygen can be transferred from the 
air to the sulphurous acid. This is accomplished by in- 
troducing the gas, sulphur dioxide, into large chambers 
together with compounds of nitrogen and oxygen, and 
steam. The compound of nitrogen and oxygen that plays 
the chief part in the transformation is the trioxide N 2 O s . 
The change is, however, not one of simple direct oxidation, 
but it involves a number of reactions. Nitric acid is used 



238 INTRODUCTION TO CHEMISTRY. 

as the starting-point. This at first reacts with sulphur 
dioxide and steam, as represented in the equation : 

2HN0 3 + 2S0 2 + H 2 = 2H 2 S0 4 + N 2 3 . 

After this the main reactions are (1) the formation of a 

compound of the formula S0 2 <qtt 2 , called nitrosyl-sul- 

phuric acid ; and (2) the decomposition of the nitrosyl- 
sulphuric acid by water. These reactions are represented 
in the two following equations : 

2S0 2 + N.O, + 0, + H 2 = 2S0 2 (OH)(N0 2 ); 

2S0 *<0H + H *° = 2S0 '<0H + N A- 

The nitrogen trioxide formed in the second reaction then 
again enters into combination with sulphur dioxide, oxygen, 
and water to form nitrosyl-sulphuric acid, which again 
undergoes decomposition with water. It will be seen, 
therefore, that the trioxide is not lost, but that it simply 
serves the purpose of effecting the oxidation of the sul- 
phur dioxide, and, theoretically, a small quantity of the gas 
is capable of transforming an infinite quantity of sulphur 
dioxide into sulphuric acid. 

Other reactions besides those mentioned above are un- 
doubtedly involved in the manufacture of sulphuric acid, 
but, according to the most elaborate researches made in 
a factory in operation, those mentioned are the principal 
ones. Whatever the details may be, the oxidation is 
effected without difficulty, and the waste in nitrogen com- 
pounds is now but slight. 

In the manufacture of sulphuric acid, sulphur is burned 
and the sulphur dioxide conducted into large chambers 
lined with lead. The reason why lead is used is that sul- 
phuric acid acts upon most other available substances. 



PROPERTIES OF SULPHURIC AGID. 239 

The Product. — The acid obtained from the leaden cham- 
bers contains about 64 per cent of sulphuric acid. It is 
evaporated in lead pans until it reaches the specific gravity 
1.75. As stronger acid acts upon lead, the evaporation is 
carried on beyond this in platinum, glass, or iron. The 
strong acid thus obtained is the concentrated sulphuric 
acid of commerce, commonly called oil of vitriol. 

Properties of Sulphuric Acid.— Sulphuric acid is an oily 
liquid, usually somewhat colored by impurities. The pure 
acid is a colorless liquid at ordinary temperatures. When 
cooled down it forms crystals. It decomposes the salts of 
most other acids, setting the acids free, and appropriating 
the metals. We have already had illustrations of this 
power in the liberation of nitric and hydrochloric acids 
from their salts by treatment with sulphuric acid. 

[Give the equations representing the action which takes 
place when common salt and potassium nitrate are treated 
with sulphuric acid.] 

Sulphuric acid unites readily with water, forming com- 
pounds with it. The simplest of these is the hydrate 
H 2 S0 4 + H 2 0. This is a crystallized substance which 
melts at a low temperature (7.5°). In consequence of the 
formation of these hydrates, a great deal of heat is evolved 
when sulphuric acid is mixed with water. This fact has 
been repeatedly illustrated in experiments already per- 
formed; and the necessity for precaution in mixing the 
two liquids has been emphasized. The acid acts upon 
organic substances containing hydrogen and oxygen, and 
extracts these elements in the proportions to form water. 
If a piece of wood is put in the acid it is charred, in con- 
sequence of the abstraction of hydrogen and oxygen. 
[How is wood usually charred in the preparation of char- 
coal ? Is there any analogy between the preparation of 
charcoal in the ordinary way and by the action of sulphuric 
acid ?] Wounds caused by sulphuric acid are painful and 
difficult to heal. 



240 INTRODUCTION TO CHEMISTRY. 

Uses of Sulphuric Acid.— Sulphuric acid is one of the 
most important chemical substances. It is used in enor- 
mous quantities in the manufacture of chlorine, of sodium 
carbonate, of artificial fertilizers, of nitroglycerin, of 
glucose, etc., and in the refining of petroleum. 

The acid is used for the purpose of drying gases upon 
which it does not act. [Can it be used for drying ammo- 
nia ?] 

Monobasic and Dibasic Acids.— Sulphuric acid differs 
markedly from nitric and hydrochloric acids in one re- 
spect. It has the power to form two different salts with 
the same metal, in one of which there is twice as much 
metal as in the other. If to a given quantity of sulphuric 
acid there is added only half the quantity of caustic potash 
required to neutralize it, a salt is formed which crystallizes. 
It has the composition represented by the formula KHS0 4 . 
If nitric acid is treated in the same way, only half the acid 
is acted on, and the product is potassium nitrate, KN0 3 , 
the rest of the acid being left unacted upon. In the case 
of sulphuric acid two reactions are possible, viz. : 

H 2 S0 4 + KOH = KHS0 4 + H 2 0, and 
H 2 S0 4 + 2KOH = K 2 S0 4 + 2H a O. 

In the case of nitric acid only one reaction seems to be 
possible: 

HN0 3 + KOH = KNO3 + H 2 0. 

Acids which, like sulphuric acid, have the power to 
form two salts with the same univalent metal are called 
dibasic acids. Acids which, like nitric acid, have the power 
to form only one salt with the same univalent metal are 
called monobasic acids. This power is connected with the 
number of replaceable hydrogen atoms contained in the 
molecule of the acids. An acid containing two replace- 



VARIOUS ACIDS OF SULPHUR. 241 

able hydrogen atoms in its molecule is dibasic; one con- 
taining one replaceable hydrogen atom in its molecule is 
monobasic. 

Acid, Neutral, and Normal Salts.— A dibasic acid yields 
two classes of salts: (1) those in which all the hydrogen is 
replaced, and (2) those in which half the hydrogen is re- 
placed by metal. The former are called normal salts, the 
latter acid salts. Normal salts are generally neutral, and 
are sometimes called neutral salts. 

Other Acids containing Sulphur.— Besides sulphurous 
and sulphuric acids, sulphur forms several other acids. 
These cannot be treated here. Their names and formulas 
are as follows : 

Hyposulphurous acid, H 2 S0 2 ; Pyrosulphuric acid, H,S 2 7 ; 
Thiosulphuric acid, H 2 S 2 3 ; Trithionic acid, H 2 S 3 6 ; 
Dithionic acid, H 2 S 2 O c ; Tetrathionic acid, H 2 S 4 6 . 

The sodium salt of thiosulphuric acid, Na,S,0 3 , com- 
monly called sodium hyposulphite, is used in photography. 
Pi/rosulphuric acid, or fuming sulphuric acid, breaks up 
into sulphuric acid and sulphur trioxide, H 9 S 2 7 = H 2 S0 4 
-f S0 3 , and is a powerful reagent for some purposes. 

Carbon Bisulphide, CS,. — Sulphur forms with carbon a 
compound known as carbon bisulphide, which has the com- 
position represented by the formula CS 2 . It is made by 
bringing carbon and sulphur together at high tempera- 
tures. It is a liquid which boils at 47°. That it dissolves 
sulphur has already been seen (see Experiments 9 and 112). 
It also dissolves many other substances. 

Selenium and Tellurium and their Compounds. — These 
elements are rarely met with. In general, their properties 
are very similar to those of sulphur, and they form com- 
pounds analogous to the principal compounds of sulphur. 
They combine with hydrogen, forming gases which have 



242 INTRODUCTION TO CHEMISTRY. 

bad odors, somewhat resembling the smell of hydrogen sul- 
phide. They burn in oxygen, forming oxides, Se0 2 and 
TeO a . Corresponding to these oxides there are acids, 
H.,Se0 3 and H 2 Te0 3 , analogous to sulphurous acid, and 
H,Se0 4 and H/IeO,, analogous to sulphuric acid. The 
compounds with hydrogen are less stable than hydrogen 
sulphide. 

The relation between the atomic weights of the three 
elements of the sulphur group has already been referred 
to. 

Points of Resemblance between Oxygen and the Members 
of the Sulphur Group. — Between the elements oxygen and 
sulphur there is very little resemblance, but the compounds 
of the two elements present many points of analogy. This 
is seen particularly in the compounds which they form 
with hydrogen and with the metals. Water and hydrogen 
sulphide are analogous in composition and in their de- 
compositions. This is also markedly true of the metallic 
oxides and sulphides; and of the hydroxides and hydro- 
sulphides. On the other hand, oxygen is unique in many 
respects, and is certainly not nearly as closely related to 
sulphur as selenium and tellurium are. 



CHAPTER XVII. 

THE NITROGEN GROUP- NITROGEN, PHOSPHORUS, AR- 
SENIC, ANTIMONY, AND BISMUTH. 

General. — Between the element nitrogen and the other 
elements which are included in this group there is but 
little resemblance. Nitrogen is a very inactive element. 
Phosphorus, on the other hand, is one of the most active. 
Nitrogen does not combine directly with oxygen. Phos- 
phorus combines with oxygen even at ordinary tempera- 
tures, while at the burning-temperature the combination 
takes place violently. The elements arsenic and antimony 
resemble each other in many respects, and are also allied to 
phosphorus; and antimony and bismuth resemble each 
other closely. On studying the compounds which all the 
members of the family form, the fact that they are closely 
related is clearly recognized. 

Phosphorus, P (At. Wt. 31). — This element occurs in 
nature in the form of phosphates, or salts of phosphoric 
acid. The chief of these is calcium phosphate, which is 
the principal constituent of the minerals phosphorite and 
apatite and of the ashes of bones. 

Preparation. — It is prepared from bone-ash. This is 
first treated with sulphuric acid. The acid converts it 
into a compound, w r hich, when mixed with charcoal and 
heated, is reduced, yielding free phosphorus. The phos- 
phorus thus obtained is cast into sticks tinder water, and 
preserved under water. 



243 



244 INTRODUCTION TO CHEMISTRY. 

Properties. — It is colorless or slightly yellow and trans- 
lucent. At ordinary temperatures it can be cut like wax, 
but it becomes hard and brittle at lower temperatures. It 
melts at a low temperature (44°) and boils at 290°. Un- 
less carefully protected from the light, its appearance 
changes. It becomes opaque and darker in color, and 
finally dark red. This change can be hastened by heating 
the phosphorus in a sealed tube to a temperature of about 
250°. 

It is insoluble in water, but soluble in carbon bisulphide. 
It is very poisonous, the inhalation of the vapor in small 
quantities causing serious disturbance of the system. The 
workmen in the factories in which phosphorus is made or 
used are frequently affected by phosphorus-poisoning. 
Among the prominent symptoms is gradual decomposition 
of the bones. 

In contact with the air phosphorus gives off fumes which 
emit a pale light visible in a dark room. It takes fire when 
rubbed or cut, and must hence be handled with great care. 
It should always be cut under water, and never held in the 
hand. It not only combines with oxygen easily, but with 
other elements, such as chlorine, bromine, and iodine, the 
action in each case being accompanied by an evolution of 
heat and light. The combination of phosphorus with oxy- 
gen has already been seen. Its conduct towards iodine can 
be shown by a very simple experiment. 

Experiment 118. — Bring together in a porcelain crucible or 
evaporating-dish a little phosphorus and iodine. It will be seen 
that simple contact is sufficient to cause the two substances to 
act upon each other. Direct combination takes place, and the 
action is accompanied by light and heat. 

Red Phosphorus. — The red substance formed when or- 
dinary phosphorus is left in the light, or heated without 
access of air, is a second variety of phosphorus known as 
red phosphorus. This differs from ordinary phosphorus 



PHOSPHOR US-PHOSPHINR 245 

as much as graphite differs from the diamond. Ordinary 
phosphorus is very active, combining readily with oxygen ; 
it is soluble in carbon bisulphide, and is poisonous. Red 
phosphorus, on the other hand, is inactive.- It does not 
change in the air, and must be heated to a comparatively 
high temperature before it will combine with oxygen; it 
is insoluble in carbon bisulphide, and is not poisonous. 
Red phosphorus is converted into the ordinary variety by 
heating it to about 300°. 

The cause of the great difference in the properties of 
the two varieties of phosphorus is not known. 

There are some other modifications of phosphorus, but 
they are rarely met with. 

Applications of Phosphorus. — Phosphorus is used princi- 
pally in the manufacture of matches, and as a poison for 
vermin. Various mixtures are used for making matches. 
Nearly all of them contain phosphorus together with some 
oxidizing compound, and some neutral substance to act as 
a medium for holding the constituents together. An ex- 
ample is a mixture consisting of 2 parts phosphorus, 1 part 
manganese dioxide, 3 parts chalk, ^ part lamp-black, and 
5 parts glue. The mixture used for "safety-matches" 
consists of potassium chlorate, minium, and antimony 
trisulphide. This will not ignite by simple friction, but 
will ignite if drawn across a paper on which is a mixture 
of red phosphorus and antimony pentasulphide. 

Phosphine, Phosphuretted Hydrogen, PH 3 . — The chief 
compound of phosphorus and hydrogen is phosphine, PH 3 . 
It is made by dissolving phosphorus in caustic potash or 
soda. The reaction which takes place is not altogether 
simple, and need not be explained at present. The points 
of chief interest in regard to the substance are : (1) its 
composition, PH 3 , which is analogous to that of ammonia, 
XH 3 ; (2) its power to combine with some acids as ammonia 



246, 



INTRODUCTION TO CHEMISTRY. 



does, forming unstable phosphonium salts analogous to the 
ammonium salts; and (3) its power to take fire when 
brought in contact with the air. 
It has a disagreeable odor. 

Experiment 119.— Arrange an apparatus as shown in Fig. 52. 
In the flask B, which should not be larger than the 100-ce. or 
125-ce. size, put about 5 grams caustic potash, dissolved in 10- 
15 cc. water, and, after the solution has become quite cold, add a 
few small pieces of phosphorus the size of a pea, and push the 




Fig. 52. 

stopper in tight. Pass hydrogen free from air for some time 
through the apparatus from the generating-flask A until all the 
air is displaced ; then disconnect at 2), leaving the rubber tube, 
closed by the pinch-cock, on the tube which enters the flask. 
Gently heat the contents of the retort, when gradually a gas will 
he evolved and will escape through the water in C. As each bub- 
ble comes in contact with the air it takes fire, and the products 
of combustion arrange themselves in rings which become larger 
as they rise. They are extremely beautiful, particularly if the , 
air of the room is quiet. Both the phosphorus and the hydrogen 
combine with oxygen in the act of burning. 

Phosphine itself does not Take Fire Spontaneously. — The 

spontaneous inflammability of phosphine has been found 



PHOSPHINE-PHOSPHORIC ACID. 24TI 

to be due to the presence of a small quantity of another 
compound of phosphorus and hydrogen which is formed by 
the action of phosphorus on caustic potash. This is a 
liquid and has the composition P 2 H 4 . It is decomposed 
by exposure to light, so that phosphine which is thus 
exposed loses the power to take fire by contact with the 
air. 

Compounds of Phosphorus with Oxygen and with Hydro- 
gen and Oxygen. — The product formed by the combina- 
tion of phosphorus and oxygen has the composition ex- 
pressed by the formula P 2 5 . This combines with water 
in different proportions, forming two distinct acids, known 
as metaphosplioric and orthopliosphoric acids : 

P 2 6 + H 2 = 2HP0 3 ; 

Metaphosphoric acid. 

P,0 6 + 3H,0 = 2H 3 P0 4 . 

Orthophosphoric acid. 

Orthophosphoric or Ordinary Phosphoric Acid, H 3 P0 4 , is 

the principal compound of phosphorus. It is the final 
product of the action of air and moisture on the element. 
As has been stated, it occurs in nature as the calcium salt 
in phosphorite and apatite. This salt is also the chief con- 
stituent of bone-ash. 

It can be made by treating bone-ash with sulphuric acid, 
and by oxidizing phosphorus. 

It is a solid crystallized substance. 

Phosphoric acid has the power of forming three distinct 
salts with the same metal. It is hence called tribasic. 
With sodium, for example, it forms the three salts Na 8 P0 4 , 
Na 2 HP0 4 , and NaH 2 P0 4 . Its normal calcium salt— that 
is to say, the one in which all the three hydrogen atoms 
are replaced by calcium — has the formula Ca 3 (P0 4 ) 2 , three 
bivalent calcium atoms replacing six atoms of hydrogen. 

[Write the equation expressing the action which takes 



248 INTRODUCTION TO CHEMISTRY. 

place when sulphuric acid decomposes normal calcium 
phosphate, forming calcium sulphate and phosphoric acid.] 

Metaphosphoric Acid, HP0 3 . — When phosphoric acid is 
heated to a sufficiently high temperature, it loses hydrogen 
and oxygen in the form of water and yields metaphosphorir 
acid : 

H 3 P0 4 = HP0 3 + H 2 0. 

Metaphosphoric acid is the substance found in com- 
merce under the name of glacial phosphoric acid. It is 
made by evaporating solutions of phosphoric acid down to 
dryness and heating the residue. 

When a salt, like ordinary sodium phosphate, HNa 2 P0 4 , 
is heated, it loses water and yields a salt of pyrophosphoric 
acid : 

2HNa 2 P0 4 = Na 4 P 2 7 + H 2 0. 

It will thus be seen that ordinary phosphoric acid by 
losing water yields pyrophosphoric acid, H 4 P 2 7 , and 
metaphosphoric acid, HP0 3 . Both these acids take up 
w r ater and are reconverted into ordinary phosphoric acid : 

HP0 3 + H 2 = H 3 P0 4 , and 
H 4 P 2 7 + H,0 = 2H 3 P0 4 . 

Phosphorous Acid, H 3 P0 3 . — This acid is formed by allow- 
ing moist air tc act on phosphorus. There is an oxide, 
P 2 3 , which bears to the acid the same relation that phos- 
phorus pentoxide bears to phosphoric acid : 

P 2 5 + 3H 2 = 2H 3 P0 4 ; 
P 2 3 + 3H.0 - 2H 3 P0 3 . 

Arsenic, As (At. Wt. 75). — Arsenic occurs in nature in 
combination with metals — as, for example, iron, copper, 
cobalt, nickel, etc. — and in combination with oxygen, as 
the trioxide As 2 3 . 

It is generally obtained by heating arsenical pyrites, 



ARSENIC— ABSINE. 



249 



FeAsS, when the arsenic separates from the iron and 
sulphur : 

FeAsS = FeS + As. 

It is also made by reducing arsenic trioxide: 

As 2 3 + 3C = 3CO + 2As. 

It has a metallic lustre. When heated to a high tern 
perature in the air it takes fire, and burns with a bluish 
flame, giving off fumes which have the odor of garlic and 
are poisonous. 

It combines directly with most elements. In the ele- 
mentary form it is not poisonous, but when oxidized it 
becomes so. 

Arsine, Arseniuretted Hydrogen, AsH 3 . — This compound 
is analogous to ammonia and phosphine. It is made by 
the action of nascent hydrogen [What is nascent hydrogen ?] 
on the compounds of arsenic with oxygen, as when these 
compounds are brought into a vessel containing zinc and 
sulphuric acid. 




Fig. 53, 



Experiment 120. — Arrange an apparatus as shown in Fig. 53. 
Put some granulated zinc in the Wolff's flask and pour dilute 



250 INTRODUCTION TO CHEMISTRY. 

sulphuric acid on it. When the air is all out of the vessel and 
the hydrogen is lighted, add slowly a little of a solution of arsenic 
trioxide, As 2 3 , in dilute hydrochloric acid. The appearance of 
the flame will soon change, becoming paler, with a slightly bluish 
tint, and giving off white fumes. (See Experiment 121.) 

Properties of Arsine.-— -Arsine is a colorless gas which is 
poisonous and has an unpleasant odor. When lighted it 
burns with a bluish-white flame, forming arsenic trioxide 
and water. It is very unstable, breaking up into arsenic and 
hydrogen when heated. When a cold object, as a piece of 
porcelain, is brought into the flame of burning arsine, the 
arsenic is deposited in the form of a dark spot. This fact 
is taken advantage of for the purpose of testing for arsenic 
in examining the stomach and other viscera in a case 
of suspected poisoning. The method is known as Marsh's 
test, as it was devised by a chemist by the name of Marsh. 

Experiment 121. — Into the flame of the burning hydrogen and 
arsine produced in the last experiment introduce a piece of porce- 
lain, as the bottom of a small porcelain dish or a crucible, and 
notice the appearance of the spots. Heat by means of a Bunsen 
burner the tube through which the gas is passing, which should 
be of hard glass. Just in front of the heated place there will be 
deposited a thin layer of metallic arsenic, commonly called a 
mirror of arsenic. This deposit is due to the direct decomposi- 
tion of the arsine into arsenic and hydrogen by heat. [Compare 
ammonia, phosphine, and arsine with reference to their stability.] 

Arsine has no basic properties, differing markedly in 
this respect from ammonia. Phosphine, as has been stated, 
has weak basic properties. 

Arsenic Trioxide, As 2 3 . — When arsenic is burned in the 
air or in oxygen it forms the trioxide. [Compare with 
phosphorus in this respect.] This substance, which is 
generally called arsenic, is made by heating compounds of 
arsenic and metals in contact with the air. Under these 



ACTD8 OF ARSENIC— ANTIMONY. 251 

circumstances, both the metal and the arsenic are oxidized, 
and the oxide of arsenic, being volatile, passes off and is 
condensed and collected in large chambers of masonwork. 

It is a colorless, amorphous, glassy mass. It is difficultly 
soluble in water, more easily in hydrochloric acid. It has 
a weak, disagreeably sweet taste, and acts very poisonously. 
It is probably more frequently used as a poison than any 
other substance. Minute quantities can be detected by the 
chemist with absolute certainty. 

The oxide is easily reduced by means of carbon. 

Experiment 122. — Mix together about equal small quantities 
of arsenic trioxide and finely-powdered charcoal. Heat the mix- 
ture in a small dry tube of hard glass, closed at one end. The 
arsenic which is set free will be deposited on the walls of the tube 
in the form of a mirror, like that obtained in Experiment 121. 

Acids of Arsenic. — Arsenic forms with oxygen and hy- 
drogen an acid of the formula H 3 As0 4 , known as arsenic 
acid, which is analogous to orthophosphoric acid. When 
heated, it undergoes changes similar to those referred to in 
connection with phosphoric acid, the products being 
metarsenic acid, HAs0 3 , and pyroarsenic acid, H 4 As 2 7 . 

When arsenic trioxide is treated with bases in solution, 
salts of arsenious acid, or the arsenites, are formed. The 
formula of the potassium salt is K 3 As0 3 . The acid 
H 3 As0 3 differs from arsenic acid, H 3 As0 4 , by one atom of 
oxygen in the molecule. 

Antimony, Sb (At. Wt. 120). — This element occurs most 
frequently in combination with sulphur as the sulphide 
Sb 2 S 3 . It is a silver-white, metallic-looking substance. 
At ordinary temperature it is not changed by contact with 
the air; but when heated to a sufficiently high temperature 
in the air it takes fire and burns, forming the white oxide. 

Experiment 123.— Heat a small piece of antimony on charcoal 
by means of the blowpipe. Notice the formation of the white 
coating on the charcoal around the place where the substance 



252 INTRODUCTION TO CHEMISTRY. 

burns. What difference is there between the conduct of anti- 
mony and arsenic before the blowpipe? See whether you can 
distinguish between antimony and arsenic by means of the blow- 
pipe. 

Stibine, Antimoniuretted Hydrogen, SbH 3 . — This com- 
pound is made in the same way as arsine. 

Experiment 124. — Make some stibine, using a solution of tartar 
emetic. 

Its properties are very much like those of arsine. It 
burns with a similar flame and is decomposed in the same 
way. 

Experiment 125. — Introduce apiece of porcelain into the flame 
and notice the deposit or antimony spot. It is darker and more 
smoky than the arsenic spot. There are other differences in 
properties by which they can be distinguished from each other 
with absolute certainty. 

Acids of Antimony. — Antimony forms acids resembling 
phosphoric, metaphosphoric, and pyrophosphoric acids. 

Antimony as a Base-forming Element. — Antimony not 
only forms acids with hydrogen and oxygen, but it also 
forms bases. These bases neutralize acids and form salts 
in which the hydrogen of the acids is replaced by antimony. 
Some of these salts are rather complicated in composition, 
and it would lead too far to discuss them here. It is suffi- 
cient for the present to recognize the important fact that 
one and the same element has the power to form acids and 
bases. 

Antimony, however, is not the only element thus far 
studied w r hich has this double power. The compounds of 
nitrogen with hydrogen and oxygen have, in general, acid 
properties, but ammonia has strongly basic properties. 
We see, therefore, that when nitrogen is combined with 
hydrogen the product has basic properties, while when 



BISMUTH. 253 

combined with hydrogen and oxygen in forms in which 
the oxygen is in excess the products are acids. The same 
is true to some extent of phosphorus. 

At the same time, neither the element nitrogen nor the 
element phosphorus itself has the power to replace the 
hydrogen of acids, and this power antimony has. 

Bismuth, Bi (At. Wt. 208). — Bismuth is not abundant 
nor widely distributed in nature. It occurs for the most 
part native, or uncombined, in veins of granite or clay 
slate. It occurs also as the oxide Bi 2 3 , and as the sul- 
phide Bi 2 S 3 . It is a hard, brittle, reddish-white substance 
with a metallic lustre. It looks very much like antimony, 
but is distinguished from it by its reddish tint. At 
ordinary temperature it remains unchanged in the air, but 
when heated to redness it burns with a bluish flame, form- 
ing the yellow oxide Bi 2 3 . 

Experiment 126.— Heat a small piece of bismuth on charcoal 
by means of the blowpipe. Compare the result with the results 
obtained with antimony and with arsenic. 

Salts of Bismuth. — Bismuth forms two classes of salts, 
known as bismuth and bismuthyl salts. In the former the 
bismuth acts as a trivalent metal, taking the place of three 
atoms of hydrogen as in the nitrate Bi(N0 3 ) 3 . In the bis- 
muthyl salts the group bismuthyl, BiO, takes the place of 
one atom of hydrogen, as in bismuthyl nitrate BiO(N0 3 ). 
Antimony forms salts similar to both these classes. 

Bismuth Nitrates. — When bismuth is dissolved in nitric 
acid and the solution evaporated to dryness the salt 
Bi(N0 3 ) 3 + 10H 2 O is obtained. This salt is decomposed 
when heated, and by water, forming basic nitrates of 
bismuth. These differ in composition according to the 
method of preparation, but all are formed by incomplete . 
neutralization of the base Bi(OH) 3 . The basic nitrate of 
bismuth, or the subnitrate, as it is called in pharmacy. 



254 INTRODUCTION TO CHEMISTRY. 

is much used in medicine as a remedy in dysentery and 
cholera. It is also used as a cosmetic. 

There are three rare elements which in their chemical 
conduct resemble the members of the nitrogen group. 
These are vanadium, colurnbium, and tantalum. It would 
be unprofitable to undertake their study at this stage. 

General Remarks on the Characteristics of the Nitrogen 
Group. — The resemblance between nitrogen and phos- 
phorus is seen particularly in the compounds ammonia and 
phosphine. Between the oxides of nitrogen and of phos- 
phorus the resemblance is not striking. There are two 
oxides of nitrogen,— the trioxide, N 2 3 , and the pentoxide, 
N 2 5 , — which in composition correspond to the two oxides 
of phosphorus, P 2 3 and P 2 5 . But while the pentoxide 
of phosphorus is the most common oxide of this element, 
the pentoxide of nitrogen is obtained with greater difficulty 
than any of the other oxides of nitrogen. There are no 
compounds of phosphorus analogous to the three principal 
oxides of nitrogen, — nitrous oxide, N 2 0; nitric oxide, NO, 
and nitrogen peroxide, N0 2 . There is no acid of phos- 
phorus corresponding to nitrous acid, HNO, , and there are 
no compounds of nitrogen analogous to phosphoric acid, 
H 3 P0 4 , and pyrophosphoric acid, H 4 P 2 0.. Nitric acid, 
HN0 3 ,and metaphosphoric acid, HP0 3 , have analogous 
compositions. 

The resemblance between phosphorus, arsenic, and anti- 
mony is much more spiking than that between nitro- 
gen and phosphorus. This resemblance has already been 
noticed in the acids formed by the three elements, and in 
their hydrogen compounds, PH 3 , AsH 3 , and SbH 3 , all of 
which are analogous to ammonia. The same resemblance 
is seen in their oxides, P 2 3 , P 2 5 , As 3 3 , As 2 5 , and 
Sb 2 3 , Sb 2 5 . Their compounds with chlorine and the 
other members of the chlorine group are also strikingly 
similar. Antimony and bismuth closely resemble each 



THE NITROGEN GROUP-BORON 255 

other in some respects. The latter forms two oxides, 
Bi 2 3 , and Bi 2 5 , and forms salts in the same way that 
antimony does, but the element has very w,eak acid prop- 
erties. The elements of the nitrogen group are trivalent in 
some compounds, as in NH 3 , PH 3 , AsH 3 , PC1 3 , AsCl 3 , etc. ; 
and quinquivalent in others, as in NH 4 C1, in which the 
nitrogen is believed to hold in combination four atoms of 
hydrogen and one atom of chlorine; in PC1 B , etc., etc. 

The atomic weights are N = 14; P = 31; As = 75; Sb = 
120; Bi = 208. These figures do not all bear simple rela- 
tions to one another, but between the atomic weights of 
phosphorus, arsenic, and antimony there exists a relation 
similar to that already noticed between the atomic weights 
of chlorine, bromine, and iodine, and sulphur, selenium, 
and tellurium. Thus, P == 31, Sb =120, and As = 75 : 

31 + 120 m „ 
— — = ,5.0. 

Again, between the atomic weights of phosphorus, anti- 
mony, and bismuth the same relation exists. Thus, P = 
31, Sb = 120, Bi = 208: 

34 + 208 

— ^ = H9.5. 

Boron, B (At. Wt. 11). — Boron may conveniently be 
studied in connection with the nitrogen group, as some of 
its properties suggest those of the members of the group. 
At the same time, it presents peculiarities which distin- 
guish it from these elements. Boron occurs in nature 
in the form of boric acid, or as salts of this acid, partic- 
ularly the sodium salt, or borax. It is prepared by treat- 
ing the oxide, B 2 3 , at a very high temperature with 
sodium or aluminium. Under proper conditions it is ob- 
tained in the form of crystals which are almost as hard as 
diamonds. 



256 INTRODUCTION TO CHEMISTRY. 

At a red heat uncrystallized boron combines with nitro- 
gen very readily. The crystallized variety can be heated 
to a high temperature in the air without changing. 
These properties distinguish boron from the members of 
the nitrogen family, all of which, with the exception of 
nitrogen, combine with oxygen. Boron combines with 
chlorine, forming the chloride BC1 3 , analogous to the 
chlorides of phosphorus and arsenic, PC1 3 and AsCl 3 . 

Boric Acid, H 3 B0 3 . — The chief compound of boron is 
boric acid. It occurs in nature in large quantities, issuing 
from the earth with water-vapor in some localities, partic- 
ularly in Tuscany. The jets of steam charged with boric 
acid, which are called suffioni, are conducted into tanks of 
water, in which the acid condenses. The solution is evap- 
orated by means of the heat of the natural steam-jets, 
and finally the acid crystallizes out. The acid is also ob- 
tained from a natural magnesium salt called boracite, and 
from borax, which is a sodium salt. 

When heated to 100°, boric acid loses water and is con- 
verted into metaboric acid, HB0 2 : 

H 3 B0 3 = HB0 2 + H 2 0. 

[What is the analogous change of phosphoric acid ?] 
The acid thus obtained is analogous to nitrous acid in 
composition. 

When heated higher, a larger proportion of water is 
given off, and an acid of the formula H 2 B 4 7 , or tetrabonc 
acid, is left behind. This is the form of boric acid from 
which borax is derived. The formula of borax is Na 2 B 4 7 
+ 10H 2 O. The relation between tetraboric acid and nor- 
mal boric acid is shown by the equation 

4H 5 B0 3 = H 2 B 4 0, + 5H 2 0. 

Heated to a still higher temperature, boric acid loses all 
of its hydrogen in the form of water, and boron trioxide, or 



BORIC ANHYDRIDE. 257 

boric anhydride, B 2 3 , is left behind. [What is the signifi- 
cance of the name boric anhydride ?] 

When a solution of borax is treated with sulphuric acid, 
boric acid is set free, and crystallizes out if the solution is 
concentrated enough. 

Experiment 127. — Make a hot solution of 30 grams crystal- 
lized borax in 120 cc. water. Add slowly 10 grams concentrated 
sulphuric acid. On cooling, the boric acid will crystallize out. 
[What evidence have you that the substance which crystallizes 
out of the solution is not borax ?] Try the solubility in alcohol 
of specimens of each. [Is there any difference?] Treat a few 
crystals of borax with about 10 cc. alcohol ; pour off the alcohol 
and set fire to it. Treat a few crystals of boric acid in the same 
way. [What difference do you observe ?] 

Boric Anhydride, B 2 3 , when heated, melts and forms a 
clear glass. This has the power to dissolve many sub- 
stances which ordinary solvents will not dissolve, and some 
of the solutions thus formed are colored. This fact is 
taken advantage of in the laboratory for the purpose of de- 
tecting the presence of those substances which form col- 
ored solutions. The method of work consists in melting a 
little boric acid or borax in a loop of platinum wire, and 
then bringing a minute particle of the substance to be ex* 
amined in contact with the glass bead thus formed. When 
heated before the blowpipe it will generally dissolve. By 
holding the bead in the oxidizing flame of the blowpipe 
the substance in solution may be oxidized, and by holding 
it in the reducing flame it may be reduced. Changes of 
color may thus be produced which will aid us in determin- 
ing what substance we have to deal with. This method 
is valuable for the purposes of analysis. 

When an alcoholic solution of boric acid is lighted, it 
burns with a green flame. The salts of boric acid do not 
color the alcohol flame. [What evidence have you had of 
the truth of this statement ?] 

Boron is trivalent in most of its compounds, as in the 
chloride, BC1 3 . 

I 
\ 



CHAPTER XVIII. 

THE CARBON GROUP : CARBON AND SILICON. 
TITANIUM — ZIRCONIUM— CERIUM — THORIUM. 

Silicon, Si (At. Wt. 28). — We have already learned how 
important a part carbon plays in animate nature. It is in- 
teresting to note that silicon, which in some respects resem- 
bles carbon from a chemical point of view, is one of the 
most important constituents of the mineral or inorganic 
parts of the earth. It occurs chiefly in the form of the 
oxide, Si0 2 , commonly called silica, or silicon dioxide; and 
in combination with oxygen and several of the common 
metals, particularly with sodium, potassium, aluminium, 
and calcium, in the form of silicates. Next to oxygen, 
silicon is the most abundant element in nature. There are 
extensive mountain-ranges consisting almost entirely of 
silicon dioxide, Si0 2 , in the form known as quartz or 
quartzite. Other ranges are made up of silicates, which are 
compounds formed by a combination of silicon dioxide and 
bases. The clay of valleys, river-beds, etc., also contains 
silicon in large quantity, while the sand found so abun- 
dantly at the sea-shore is mostly silicon dioxide, Si0 2 . 

Silicon is never found in the free state, and it is difficult 
to decompose the oxide, Si0 2 , in such a way as to get the 
element, though it can be accomplished by heating the oxide 
with potassium and with magnesium. Under proper con- 
ditions silicon can be obtained in the form of crystals which 
have a gray color and are harder than glass. It is not 
acted upon by the strongest acids, nor when heated in a 
current of oxygen, 

258 



SILICIC ACID. 259 

With hydrogen, silicon forms a gaseous compound of the 
formula SiH 4 ; it combines with chlorine, forming SiCl 4 , 
and with fluorine, forming SiF 4 . The fluoride has already 
been referred to in connection with the action of hydro- 
fluoric acid on silicates. We have seen that hydrofluoric 
acid dissolves silicates — as, for example, glass — in conse- 
quence of the action of the acid on silicon dioxide, which 
is represented thus : 

Si0 2 + 4HF = SiF 4 + 2H 2 0. 

The silicon fluoride passes off in the form of gas. 

Silicic Acid. — There are several varieties of silicic acid, 
all of which are, however, derived from an acid of the 
formula H 4 Si0 4 , or normal silicic acid. When this is set 
free from its salts, it loses water, and is changed to ordinary 
silicic acid, H 2 Si0 3 : 

H 4 Si0 4 = H 2 Si0 3 + H 2 0. 

When heated, this second form of silicic acid is converted 
into the dioxide Si0 2 : 

H 2 Si0 3 = Si0 2 + H 2 0. 

Most of the ordinary silicates are derived from the acid of 
the formula H 2 Si0 3 . [What is the formula of carbonic acid ? 
Under what circumstances does carbonic acid break up 
into carbon dioxide and water?] Other silicic acids are 
obtained by heating ordinary silicic acid. Thus, under the 
proper conditions an acid of the formula H 2 Si 2 5 , and one 
of the formula ET Si,O s , are obtained: 



'«> 



2H,SiO s = H a Si,0 6 + H a O; 

3H s Si0 3 = H 4 Si 3 8 + H,0. 

These are called poly silicic acids. Some of these are found 
in mature. Opal is the best-known example. 



260 INTRODUCTION TO CHEMISTRY. 

Silicon Dioxide, Silicic Anhydride, Si0 2 . — As already 
stated, this substance occurs very abundantly in nature and 
in many "different forms. Quartz, or rock crystal, is pure 
crystallized silicon dioxide; quartzite is a coarser-grained 
substance made up of small crystals of quartz, usually col- 
ored. Agate, amethyst, and carnelian are varieties of quartz 
colored by foreign substances. 

Silicon dioxide is insoluble in water and acids. It is solu- 
ble in hydrofluoric acid, as has been stated. Glass consists 
of salts of silicic acid, usually of the sodium or potassium 
salts and calcium salts. 

Comparison of Carbon and Silicon. — The two elements of 
this family resemble each other in the composition of some 
of their simplest compounds, as carbon dioxide, C0 2 , and 
silicon dioxide, Si0 2 ; carbonic acid, H 2 C0 3 , and silicic acid, 
H 2 Si0 3 ; marsh-gas, CH 4 , and silicon hydride, SiH 4 ; carbon 
tetrachloride, C01 4 , and silicon tetrachloride, SiCl 4 . On 
the other hand, they present marked points of difference. 
Each, yields a large number of derivatives, but the deriva- 
tives of carbon bear to the element relations entirely differ- 
ent from those which the derivatives of silicon bear to this 
element. The compounds of carbon can all be shown to 
be derived from the hydrocarbons; that is to say, they may 
be regarded as formed from the hydrocarbons by a com- 
paratively simple set of changes, while most of the com- 
pounds containing silicon are derivatives of silicic acid. 

Rare Elements of this Group. — The rare elements tita- 
nium, zirconium, cerium, and thorium resemble silicon in 
their chemical conduct. They form oxides of the formulas 
Ti0 2 , Zr0 2 , Ce0 2 , and Th0 2 which are analogous to silicon 
dioxide. Titanium occurs in nature principally as the 
dioxide, and forms the three minerals rutile, brookite, and 
anatase. Zirconium occurs in nature as zircon, which is a 
silicate of the formula ZrSi(X. 



CHAPTER XIX. 
BASE-FORMING ELEMENTS.— GENERAL CONSIDERATIONS. 

Introductory. — At the end of Chapter XIV is this sen- 
tence : " After the acid-forming elements have been studied, 
the base-forming elements will be taken up in a similar 
way; but, as will be seen, the chemistry of the acid-form- 
ing elements exhibits more variety, and is hence better 
adapted to the illustration of the general principles of the 
science than that of the base-forming elements, so that the 
latter need not be treated as fully." 

The significance of the name base-forming elements has 
been stated. It is simply this: that the compounds of 
these elements with hydrogen and oxygen are bases, or, in 
other words, have the power to neutralize acids and form 
salts. But the distinction between acid-forming and base- 
forming elements is not a sharp one, for the reason that 
there are some elements that occupy an intermediate posi- 
tion, forming both acids and bases. One example of this 
kind already studied is antimony, and the reason why it 
was considered as a member of the nitrogen group is that 
it is unquestionably closely related to arsenic, which is an 
acid-forming element. A close study will show that those 
elements which have the power to form both acids and bases 
are related to one of the four groups already studied. 
There are, thus, certain elements which show some resem- 
blance to the members of the chlorine group, but never- 
theless act principally as base-formers; so, too, there are 
certain elements w r hich resemble the members of the sulphur 
group, but which generally form bases. In a similar way, 

261 



262 IMTUObVCTlON TO CHEMISTRY. 

there are base-forming analogues of the nitrogen and carbon 
groups. Those elements which always act as base-formers 
have no analogues among the acid-forming elements. 

Order in which the Base-forming Elements will be Taken 

up. — In studying the base-forming elements, it appears best 
to begin with those which have the most strongly marked 
character. These are members of Group I, as shown in the 
table page 212. It further appears best to adhere as closely 
as possible to the arrangement in the periodic system. 
Accordingly the following order will be observed in the 
presentation of the elements yet to be studied: 

1. The Potassium Group, consisting of lithium, sodium, 
potassium, rubidium, and caesium. 

2. The Calcium Group, consisting of glucinum, calcium, 
barium, and strontium. 

3. The Magnesium Group, consisting of magnesium, 
zinc, and cadmium. 

4. The Silver Group, consisting of silver, copper, and 
mercury. 

5. The Aluminium Group, of which aluminium is the 
only well-known member. Allied to it are the rare ele- 
ments gallium, indium, thallium, scandium, yttrium, lan- 
thanum, and ytterbium. 

6. The Lead Group, consisting of germanium, tin, and 
lead. 

7. The Chromium Group, consisting of chromium, mo- 
lybdenum, and tungsten. The members of this group 
show some analogy to the members of the sulphur group, 
as will be pointed out when chromium is considered. 

8. The Manganese Group, of which manganese is the 
only representative. There are some points of resemblance 
between manganese and the members of the chlorine group. 

9. The Iron Group, consisting of iron, cobalt, and nickel. 

10. The Palladium Group, consisting of palladium, 
ruthenium, and rhodium. 



METALLIC PROPERTIES. 263 

11. The Platinum Group, consisting of osmium, irid- 
ium, platinum, and gold. 

It will be seen at once that there are many more base- 
forming than acid-forming elements, and it is a serious 
undertaking to become thoroughly acquainted with all the 
elements included under this head. In order to get a gen- 
eral knowledge of the principles of chemistry, however, it 
is not necessary to study all these elements* The chemist 
must, of course, familiarize Rimself to some extent with 
all of them, and those who continue the study of chemistry 
hereafter will have abundant opportunity to study them in 
detail. For the present it will be best to confine attention 
to a few of the representative elements included in the 
above list. A knowledge of these will make it possible to 
study the others without serious difficulty, should occasion 
demand. 

Metallic Properties. — It has long been customary to 
divide the chemical elements into two classes, — the metals 
and the non-metals. This classification was originally based 
upon differences in the physical properties of the elements, 
the name metal being applied to those elements which have 
what is known as a metallic lustre, are opaque, and are 
good conductors of heat and electricity. All those elements 
which do not have these properties were called non-metals. 
Gradually the name metal came to signify an element which 
has the power to replace the hydrogen of acids and form 
salts, and the name non-metal to signify an element which 
has not this power. This classification, as will be seen, is 
about the same as that made use of in this book. It thus 
appears that, in general, elements that have similar physi- 
cal properties have also similar chemical properties. 

Occurrence of the Metals. — One of the first questions that 
suggests itself in connection with each element is, in what 
forms of combination does it occur in nature ? The chem- 



264 INTRODUCTION TO CHEMISTRY. 

ical elements and compounds that occur ready-formed in 
nature are called minerals; and the minerals and mixtures 
of minerals from which the metals are extracted for prac- 
tical purposes are called ores. The most common ores are 
oxides and sulphides. Examples of these are the ores of 
iron, tin, copper, lead, and zinc. The carbonates also 
occur in large quantity in nature, and are used for the 
purpose of preparing some of the metals. The carbon- 
ate of zinc, for example, is a'valuable ore of zinc. 

Extraction of Metals from their Ores. — The detailed 
study of the methods used in the extraction of the metals 
from their ores is the object of metallurgy. Besides the 
methods used on the large scale, there are others which are 
only used in the laboratory. The most common method of 
extracting metals from their ores is that used in the case of 
iron, which consists in heating the oxides with charcoal or 
coke. If the ores used are not oxides they must first be 
converted into oxides before this method is applicable. 
This can generally be accomplished by heating the ores in 
contact with the air. Under these circumstances the nat- 
ural carbonates, sulphides, and hydroxides are converted 
into oxides. These changes are illustrated by the following 
equations : 



FeC0 3 = FeO + C0 2 ; 
2FeO + = Fe 2 3 ; 
2FeS 2 + llO = Fe 2 O s + 4S0 2 ; 
2Fe(OH) 3 = Fe 2 3 + 3H 2 0. 



A second method consists in reducing the oxide by heat- 
ing it in a current of hydrogen. This has been illustrated 
in the action of hydrogen upon copper oxide, when the 
following reaction takes place : 

CuO + H 8 = H 2 + Cu. 



COMPOUNDS OF THE METALS. 265 

The method is efficient for many oxides, but is expensive 
and is not used on the large scale. 

Another method of extraction consists in treating the 
chloride of a metal with sodium. This is illustrated in the 
preparation of magnesium: 

MgCl 2 + 2Xa = Mg + 2NaCl. 

Such a method is employed only in case it is impossible or 
extremely difficult to reduce the oxide. 

Besides the above methods, there are others which will 
be described under the individual metals. 

The Properties of the Metals. — As we shall find, the 
metals differ very markedly from one another. Some are 
light, floating on water, as lithium, sodium, etc.; some are 
extremely heavy, as lead, platinum, etc. Some combine 
with oxygen with great energy; others form very unstable 
compounds with oxygen. Some form strong bases; others 
form w T eak bases. In general, those elements which are 
lightest, or which have the lowest specific gravity, are the 
most active chemically, while those which have the highest 
specific gravity are the least active. Among the former 
are lithium, sodium, and potassium; among the latter are 
lead, gold, and platinum. 

Compounds of the Metals. — The compounds of the metals 
may be conveniently classified as: 

a. Compounds with fluorine, chlorine, bromine, and io- 
dine; or the fluorides, chlorides, bromides, and iodides. 

b. Compounds with oxygen, and with oxygen and hydro- 
gen; or the oxides and hydroxides. 

c. Compounds with sulphur, and with sulphur and hydro- 
gen; or the sulphides and hydrosulphides. 

d. Compounds with the acids of nitrogen; or the ni- 
trates and nitrites. 

e. Compounds with the acids of chlorine, bromine, and 



266 INTRODUCTION TO CHEMISTRY. 

iodine; or the chlorates, bromates, iodates, hypochlorites, 
etc. 

/. Compounds with the acids of sulphur; or the sul- 
phates, sulphites, etc. 

g. Compounds with carbonic acid; or the carbonates. 

h. Compounds with the acids of phosphorus, arsenic, and 
antimony; or the phosphates, arsenates, etc. 

i. Compounds with silicic acid; or the silicates. 

j. Compounds with boric acid ; or the borates. 

Of the almost infinite number of compounds belonging 
to the classes above referred to, only a comparatively small 
number will be treated of in this book. It is more impor- 
tant to become acquainted with the general methods of 
preparation and the general properties of these compounds 
than to learn details concerning many individual members 
of each class. Only those compounds will be treated which 
illustrate general principles, or which, owing to some appli- 
cation, happen to be of special interest. 

The acids of which the salts are derivatives are already 
known to us, and in dealing with acids frequent reference 
has been made to the methods of making the salts, and to 
some of their most important properties. It will be well, 
before taking up the metals systematically, to discuss 
briefly the general methods of preparation, and the general 
properties of the different classes of metallic compounds. 
It must be borne in mind, however, that the only way to 
become familiar with these substances and their relations 
is by working with them in the laboratory. 

Chlorides. — The chlorides, as well as the fluorides, bro- 
mides, and iodides, may be regarded as the salts of hydro- 
chloric, hydrofluoric, hydrobromic, and hydriodic acids, or 
simply as compounds of the metals with the members of 
the chlorine group. The most important of these com- 
pounds are the chlorides, and these well illustrate the con- 
duct of the others. 



j 



CHLORIDES, 267 

The chlorides are made by treating the metals with 
hydrochloric acid or chlorine; by treating an oxide or a 
hydroxide with hydrochloric acid; by treating an oxide 
with chlorine and a reducing agent, like carbon; by treat- 
ing a salt of a volatile acid, a carbonate for example, with 
hydrochloric acid; by treating a salt of an insoluble acid 
with hydrochloric acid; by adding hydrochloric acid or a 
soluble chloride to a solution containing a metal with 
which chlorine forms an insoluble compound; and by 
adding to a solution of a chloride a salt the acid of which 
forms with the metal of the chloride an insoluble salt, while 
the metal contained in it forms with chlorine a soluble 
chloride. 

Only two of the above methods are peculiar to chlorides. 
These are the treatment of the metals with chlorine, and 
the treatment of oxides with chlorine and a reducing agent. 
The others involve principles which are also involved in the 
preparation of all salts, and they may therefore be treated 
in a general way. 

Examples.— Zinc chloride is formed by treating zinc 
w T ith hydrochloric acid. 

[Write the equation.] 

Ferric chloride is formed by treating iron with chlorine : 

Fe + 3C1 = FeCl 3 . 

Calcium chloride is formed by treating lime or calcium 
oxide, CaO, with hydrochloric acid: 

CaO + 2HC1 = CaCl 2 + H 2 0. 

Sodium chloride is formed by treating sodium hydroxide, 
or caustic soda, NaOH, with hydrochloric acid: 

NaOH + HG1 == NaOl + H 2 0. 

[What takes place when caustic soda or caustic potash is 
treated with chlorine ?] 



268 INTRO D UCTION TO CHEMISTR Y. 

Calcium chloride is formed when calcium carbonate, 
CaC0 3 , is treated with hydrochloric acid: 

CaCO, + 2HC1 = CaCl 2 + C0- a + H 2 0. 

Silver chloride is precipitated when hydrochloric acid or 
a soluble chloride is added to a solution containing a silver 
salt. 

Experiment 128. — Dissolve a small crystal of silver nitrate in 
pure water. Add to a small quantity of this solution in a test-tube 
a few drops of dilute hydrochloric acid. The white substance thus 
precipitated is silver chloride, AgCl. To another small portion 
of the solution add a few drops of a dilute solution of common 
salt, or sodium chloride, NaCl. The white substance produced 
in this case is also silver chloride. Add ammonia to each tube. 
If sufficient is added the precipitates will dissolve. On adding 
enough hydrochloric acid to these solutions to combine with all 
the ammonia, the silver chloride is again thrown down. On 
stauding exposed to the light both precipitates change color, 
becoming finally dark violet. The reactions involved in the above 
experiments are these: In the first place, when hydrochloric acid 
is added to silver nitrate this reaction takes place: 

AgN0 3 + HC1 = AgCl + HN0 3 . 

When sodium chloride is added this reaction takes place: 

AgNOs + NaCl = AgCl + NaN0 3 . 

In the first reaction nitric acid is set free; in the second, the 
sodium and silver exchange places. In addition to the insoluble 
silver chloride, there is formed at the same time the soluble salt, 
sodium nitrate. On adding ammonia the silver chloride forms 
with it a compound w r hich is soluble in water ; and, on adding an 
acid, the ammonia combines with it, leaving the silver chloride 
uncombined and therefore insoluble. 

General Properties of the Chlorides. — Most of the chlo- 
rides of the metals are soluble in water without decom- 
position, though many of them are decomposed when heated 
to a sufficiently high temperature with water. Silver 






OXIDES— HYDROXIDES. 269 

chloride, AgCl, and mercurous chloride, HgCl, are insoluble 
in water. Lead chloride, PbCl 2 , is difficultly soluble in 
water. If, therefore, on adding hydrochloric acid or a 
soluble chloride to a solution, a precipitate is formed, the 
conclusion is generally justified that one or more of the 
three metals — silver, lead, or mercury — is present. By 
taking into account the differences between these chlorides, 
it is not difficult to decide of which of them a precipitate 
consists. 

Oxides. — These occur very generally in nature, and are 
among the most common ores of some of the important 
metals. The oxides of iron, tin, manganese, etc., are all 
found in nature. They can be made by oxidizing the 
metals, by heating nitrates and carbonates, and by heating 
hydroxides. 

Examples. — When magnesium is burned (see Experi- 
ment 14) it is converted into magnesium oxide : 

Mg + = MgO. 

When lead nitrate is heated it gives off oxygen and an 
oxide of nitrogen and leaves behind lead oxide : 

Pb(N0 3 ) 2 = PbO + 2X0, + 0. 

When calcium carbonate is heated it gives off carbon 
dioxide and leaves behind calcium oxide, CaO : 

CaC0 3 = CaO + CO,. 

Hydroxides. — The hydroxides are formed by treating 
oxides with water, and by decomposing salts by adding 
soluble hydroxides. 

Examples. — When calcium oxide or lime is treated with 
water it is converted into the hydroxide, Ca(OH) 2 , or slaked 
lime. 



270 INTRODUCTION TO CHEMISTRY. 

Experiment 129. — To some pieces of freshly-burned lime add 
enough cold water to cover it. The action which takes place is 
represented by the equation 

CaO + H 2 = Ca(OH) 2 . 

The process is known as slaking. [What evidence have you that 
heat is evolved in the reaction, and that the substance obtained 
is not calcium oxide ?] 

Most of the hydroxides of the metals are insoluble in 
water. If a soluble hydroxide is added to a solution con- 
taining a metal whose hydroxide is insoluble, the latter is 
precipitated. Thus, if a solution of sodium hydroxide is 
added to a solution of a magnesium salt, magnesium hy- 
droxide is precipitated : 

MgS0 4 + 2NaOH = Na 2 S0 4 + Mg(OH) 2 . 

Experiment 130. — To a small quantity of a dilute solution of 
magnesium sulphate add a dilute solution of caustic soda. The 
white precipitate is magnesium hydroxide. [Would you expect 
this precipitate to be soluble in sulphuric acid ? in hydrochloric 
acid ? in nitric acid ?] The answers follow from these considera- 
tions: When acids act upon hydroxides, salts are formed; mag- 
nesium sulphate is soluble, as is seen by the fact that we started 
with a solution of this salt ; the only insoluble chlorides are those 
of silver, lead, and mercury; all nitrates are soluble. 

When a solution of an iron salt is treated with sodium 
hydroxide a precipitate of iron hydroxide is formed : 

FeCl 3 + 3NaOH = Fe0 3 H 3 + 3NaCl. 

Experiment 131. — To a dilute solution of that chloride of 
iron which is known as ferric chloride add caustic soda. The 
reddish precipitate formed is ferric hydroxide. [From the gen- 
eral statements made above, would you expect this precipitate to 
be soluble in sulphuric acid? in hydrochloric acid? in nitric 
acid? Try each.] 



DECOMPOSITION OF SALTS. 271 

Only the hydroxides of the members of the potassium 
family and of the calcium family are soluble in water. 
The hydroxides of sodium and potassium are called alka- 
lies. The solution of ammonia in water acts like a soluble 
hydroxide and probably contains ammonium hydroxide, 
NH 4 OH, formed by the action of water on ammonia: 

NH 3 + H 2 = NH 4 OH. 

When any one of the soluble hydroxides is added to a salt 
containing any metal that does not belong to the potassium 
or calcium family, an insoluble compound is thrown down. 

[Test this by trying such salts as may be available. 
Note the results in each case. Is an insoluble compound 
formed ? What is its general appearance ?] 

Decomposition of Salts by Acids and by Bases. — The de- 
composition of salts by the addition of hydroxides is^in 
some respects analogous to the decomposition of salts by 
the addition of strong acids. 

When an acid is added to a salt there are three cases 
which may present themselves: 

1. The acid from which the salt is derived may be vola- 
tile or may break up, yielding a volatile product. 

In this case decomposition takes place, and the volatile 
acid is given off. This is illustrated by the liberation of 
hydrochloric and nitric acids from chlorides and nitrates 
by the addition of sulphuric acid, and of carbon dioxide 
from carbonates by the addition of other acids. 

[Write the equations representing the action which 
takes place when sulphuric acid acts upon potassium chlo- 
ride, calcium chloride, sodium nitrate, calcium nitrate; 
when hydrochloric acid acts upon sodium carbonate, cal- 
cium carbonate.] 

2. The acid from which the salt is derived may be in- 
soluble or difficultly soluble in water, and not volatile. 

In this case, if the salt is in solution, decomposition takes 



272 INTRODUCTION TO CHEMISTRY. 

place, and the insoluble or difficultly soluble acid is precip- 
itated. This is illustrated by the liberation of boric acid 
from borax by the addition of sulphuric acid; and by the 
liberation of silicic acid by the addition of hydrochloric or 
sulphuric acid to a soluble silicate. 

3. The acid from which the salt is derived may be 
soluble and not volatile under the existing conditions. 

In this case, if the substances are in solution, apparently 
no change takes place. Thus, when nitric acid is added to 
sodium chloride in solution no striking change takes place, 
no gas is given off, no precipitate is formed. It is ex- 
tremely difficult to determine what does take place under 
these circumstances. A study of such cases as this is of 
great importance to chemistry, but cannot be undertaken 
at this stage. 

Now, to return to the action of hydroxides upon salts; 
when a soluble base acts upon a salt, three cases may pre- 
sent themselves: 

1. The base from which the salt is derived may be vola- 
tile or may break up, yielding a volatile product. 

In this case decomposition takes place and the volatile 
base is given off. This is not a common case except among 
the compounds of carbon. The one illustration which we 
have had is the decomposition of ammonium salts by cal- 
cium hydroxide and by sodium hydroxide. 

[Write the equations representing the action in both 
oases. In what does the analogy between the decomposi- 
tion of ammonium salts by bases and of carbonates by 
acids consist ?] 

2. The hydroxide or base from wiiich the salt is derived 
may be insoluble or difficultly soluble in water, and non- 
volatile. 

In this case, if both the salt and the base are in solution, 
decomposition takes place, and the insoluble or difficultly 
soluble hydroxide or base is precipitated. This has already 
been illustrated. 



SULPHIDES. 273 

3. The base from which the salt is derived may be solu- 
ble and not volatile. 

In this case there is no direct evidence of change. 
Thus, when sodium hydroxide is added to potassium ni- 
trate, nothing is seen except a clear solution. To deter- 
mine what takes place is a difficult matter.* 

Sulphides. — Many sulphides are found in nature. They 
are made by heating metals with sulphur; by treating so- 
lutions of salts with hydrogen sulphide or soluble sul- 
phides. 

Examples. — Among the common natural sulphides are 
iron pyrites, FeS, ; lead sulphide, or galenite, PbS; copper 
pyrites, FeCuS 2 . [Examine several specimens of each, and 
note their general properties.] 

When copper or iron is heated with sulphur the corre- 
sponding sulphides are formed. (See Experiments 10 and 
113.) [For what purpose were these experiments per- 
formed ?] 

When hydrogen sulphide is passed through a solution 
containing a metal whose sulphide is insoluble, the sul- 
phide is precipitated. This has been illustrated by passing 
the gas through solutions of lead nitrate, zinc sulphate, 
and arsenic trioxide. The reactions are: 

*Here a word of warning to students. Do not forget that when- 
ever a precipitate is formed there is something in the solution which 
is just as important as the precipitate. Accustom yourselves to re- 
gard every case of chemical action as a whole. The statement that 
a precipitate is formed when sodium hydroxide is added to a solution 
of an iron salt is a very imperfect description of the chemical change 
that takes place. Precipitates have come to be regarded in a false 
light, in consequence of the constant use made of them for purposes 
of analysis It must be remembered that analysis is not chemistry, 
though it is essential to the study of chemistry and is an important 
application of the science. The art of analysis is founded upon a 
knowledge of the science of chemistry. If you have a knowledge of 
the science, it will comparatively be easy to acquire the art of analy- 
sis, should this acquisition become desirable. 



274 INTRODUCTION 10 CHEMISTRY. 

Pb(N0 3 ) 2 + H 2 S = PbS + 2HN0,; 
ZnS0 4 + H 2 S = ZnS + H 2 S0 4 ; 

As 2 3 + 3H 2 S = As 2 S 3 + 3H 2 0. 

[What differences were observed in these three cases ? 
Repeat the experiments.] 

When a soluble sulphide, as ammonium sulphide or so- 
dium sulphide, is added to a solution containing a metal 
whose sulphide is insoluble, the insoluble sulphide is 
thrown down. 

Experiment 132. — Add ammonium sulphide successively to 
dilute solutions of an iron salt, a lead salt, a copper salt. Note 
what takes place in each case. 

Qualitative Analysis. — The sulphides of all the metals 
except those which belong to the potassium and calcium 
groups, and that of magnesium, are insoluble in water. 
Of those sulphides which are insoluble in water, some are 
insoluble and some are soluble in dilute hydrochloric acid. 
Further, of those which are insoluble in dilute hydrochloric 
acid, some are soluble and some are insoluble in ammonium 
sulphide. 

These facts furnish the basis of the method commonly 
employed in analyzing substances. Suppose we have a 
solution containing all the more common elements, and 
we wish to determine what is in it. Let us suppose 
that, on adding hydrochloric acid to it, a precipitate is 
formed. [What does this show ?] This precipitate is fil- 
tered off, and the solution treated with hydrogen sulphide. 
Those metals whose sulphides are insoluble in dilute hy- 
drochloric acid will be precipitated. Among the elements 
which may be contained in this precipitate are lead, mer- 
cury, copper, tin, arsenic. The solution from which the 
precipitate was thrown down may still contain those met- 
als whose sulphides are soluble in dilute hydrochloric 
acid. If, therefore, we filter off the precipitate and add 



QUALITATIVE ANALYSIS— NITRATES. 275 

ammonium sulphide to the filtrate, the metals whose sul- 
phides are insoluble in neutral or alkaline solutions will be 
thrown down. Among these are iron, aluminium, chro- 
mium, manganese, etc. The filtrate from this precipitate 
may contain all those metals whose sulphides are soluble 
in water. By means of other reactions these can be sub- 
divided into groups. In the ordinary method of analysis 
we have, therefore, several groups of elements to deal with. 
These are : 

1. The liydrochlomc-acid group, consisting of those met- 
als whose chlorides are insoluble in water. 

2. The hydrogen-sulphide group, consisting of those 
metals whose sulphides are insoluble in dilute hydrochloric 
acid. 

3. The ammonium -sulphide group, consisting of those 
metals whose sulphides are soluble in dilute hydrochloric 
acid, but are precipitated by ammonium sulphide. 

4. Elements whose sulphides are soluble in water. 

Each of these groups can be subdivided, and the sub- 
groups again subdivided, until positive evidence of the 
presence of certain metals is obtained. 

Hydrosulphides are formed when hydrogen sulphide is 
passed into a solution of a hydroxide until no more is taken 
up. 

Potassium hydrosulphide is formed thus : 

KOH + HJ8 = KSH + H 2 0. 

Ammonium hydrosulphide is formed thus: 

NH 4 OH + H 2 S - NH 4 SH + H 2 0. 

Nitrates. — These salts are formed by treating metals 
with nitric acid; by treating oxides or hydroxides with 
nitric acid, and in general by treating any easily decom- 
posed salt as a carbonate with nitric acid. 



276 INTRODUCTION TO CHEMISTRY. 

Examples. — When nitric acid acts upon copper, copper 
nitrate is formed. [What else is formed ? Give an account 
of the changes which take place. Write the equation rep- 
resenting the reaction.] 

The simple neutralization of nitric acid with a base or 
hydroxide has been illustrated in the experiments on acids 
and bases (Experiment 62). [Write the equations repre- 
senting the reactions which take place when nitric acid is 
neutralized with potassium hydroxide, with calcium oxide, 
with calcium hydroxide.] 

All nitrates are soluble in water, and all are decomposed 
by heat. [Try the solubility, in water, of such nitrates as 
may be available.] 

Experiment 133. — Heat 2 to 3 grams potassium nitrate on 
charcoal with the blowpipe flame. The decomposition with evo- 
lution of gas is called deflagration. Heat some copper nitrate 
and lead nitrate. Carefully note the changes which take place. 
The compounds left behind are copper oxide and lead oxide. 

Chlorates are made from potassium chlorate, which is 
made by treating a strong solution of caustic potash with 
chlorine. [Explain the reaction.] 

Chlorates are soluble in water, and are decomposed by 
heat with evolution of oxygen. [When potassium chlorate 
is heated, what takes place in the first stage of the opera- 
tion ?] 

Hypochlorites are formed by treating some of the metal- 
lic hydroxides in dilute solution with chlorine. This has 
been illustrated in the formation of " bleaching-powder," 
which contains calcium hypochlorite. [Explain what takes 
place when slaked lime is treated with chlorine.] 

Hypochlorites are decomposed by heat. 

Sulphates. — Some sulphates, as those of calcium and ba- 
rium, are found in nature, the former being known as gyp- 
sum. Sulphates are made by treating metals or metallic 






SULPHATES. 211 

hydroxides or oxides with sulphuric acid; by treating easily- 
decomposed salts, as carbonates, with sulphuric acid; and 
by treating a solution containing a metal whose sulphate is 
insoluble, with sulphuric acid or a soluble sulphate. 

Examples. — Usually, when sulphuric acid acts upon a 
metal, hydrogen is evolved and a salt is formed. This has 
been illustrated in the preparation of hydrogen by means 
of zinc and sulphuric acid. 



t 



Experiment 134. — Dissolve some iron in dilute sulphuric acid. 
When the acid is neutralized, filter the solution and evaporate it 
down to crystallization. [What is the appearance of the salt ? 
Does it contain water of crystallization ? Was hydrogen evolved 
during the action of the acid on the metal ?] Dry some of the 
salt, and put it aside for further use. 

Experiment 135. — Dissolve some copper foil in concentrated 
sulphuric acid. [In what respect does the action in this case differ 
from that in the last experiment?] When the action is over, and 
the mass cooled down, pour it into three or four times its volume 
of water, when most of the black deposit will dissolve. Evapo- 
rate the solution, and get out some of the salt in the form of 
crystals. [What is the appearance of the salt ? Does it contain 
water of crystallization ? What does the salt look like after it 
has been heated in a tube ?] Dry some of it, and put it aside for 
further use. [Write the equations representing the action which 
takes place when copper acts upon sulphuric acid.] 

The action of sulphuric acid on metallic hydroxides has 
been illustrated. (See p. 119.) 

[Write the equation representing the action which takes 
place when the acid acts upon sodium hydroxide, potassium 
hydroxide, ammonium hydroxide. AVhat is monosodium 
sulphate ? What is neutral sodium sulphate ? Is there 
any difference between disodium sulphate and neutral so- 
dium sulphate ?] 

Most sulphates are soluble in water. The sulphates of 
barium, strontium, and lead are insoluble in water, and the 
sulphate of calcium is difficultly soluble. Therefore, when 



278 INTRODUCTION TO CHEMISTRY. 

sulphuric acid is added to a solution containing either 
of the metals barium, strontium, or lead, a precipitate is 
formed. 

Experiment 136. — Make a dilute solution of barium chloride, 
of lead nitrate, of strontium nitrate. To a small quantity of 
each in a test-tube add a little sulphuric acid. In each case a 
white precipitate is formed. [What remains in solution ?] Make 
a somewhat concentrated solution of calcium chloride. To this 
add some sulphuric acid. A precipitate is formed. [What is in 
solution ?] Add more water, and see whether this precipitate 
will dissolve. The formulas of the salts used in the experiments 
are barium chloride, BaCl 2 ; lead nitrate, Pb(N0 3 )2 ; strontium 
nitrate, Sr(N0 3 ) 2 . [Write the equations expressing the reactions.] 
If to the solutions of the salts any soluble sulphate is added in- 
stead of sulphuric acid, the same insoluble sulphates will be 
formed. The sulphates of iron, copper, sodium, and potassium 
are among the soluble sulphates. Make dilute solutions of small 
quantities of each of these, and add them successively to the so- 
lutions of barium chloride, lead nitrate, and strontium nitrate. 
The formula of iron sulphate is FeSC>4 ; of copper sulphate, 
C11SO4 ; of sodium sulphate, Na 2 S04 ; and of potassium sulphate, 
K2SO4. Write the equations representing the reactions which 
take place in the above experiments. It need hardly be explained 
that the action consists in an exchange of places on the part of 
the metals. Thus, when the soluble salt, iron sulphate, FeS0 4 , is 
brought together with the soluble salt, barium chloride, BaCU , the 
insoluble salt, barium sulphate, BaS0 4 , and the soluble salt, iron 
chloride, FeCl 2 , are formed: 

FeS0 4 + BaCl 2 = FeCl 2 4- BaS0 4 . 

When heated with charcoal in the reducing flame of the 
blowpipe, sulphates are reduced to sulphides: 

K,S0 4 + 4C = K 2 S + 4CO, or 
K 2 S0 4 + 2C - K 2 S + 2C0 2 . 

Experiment 137. — Mix and moisten a little sodium sulphate and 
finely-powdered charcoal. Heat the mixture for some time in the 



S U LP HUES— CA RBONA TES. 279 

reducing flame. After cooling scrape off the salt, dissolve it in a 
few cubic centimetres of water, and filter through a small filter. If 
the change to the sulphide has taken place, sodium sulphide, Na a S, 
is in solution. A soluble sulphide when added to a solution con- 
taining copper gives a black precipitate of copper sulphide. Try 
this; also try the action on copper sulphate solution of some of 
the sulphate from which the sulphide was made. 

Sulphites are made from sodium or potassium sulphite, 
which are made by treating sodium or potassium hydrox- 
ide in solution with sulphur dioxide: 

2NaOH + SO a =? Na 2 S0 3 + H 3 0. 

All sulphites are decomposed by the common acids, sul- 
phur dioxide being given off: 

Na a S0 3 + H 2 S0 4 = Na 2 S0 4 + H 2 + S0 2 . 

Carbonates. — Many carbonates are found in nature, 
some of them in great abundance, and widely distributed. 
The principal one is calcium carbonate. They are made 
by passing carbon dioxide into solutions of hydroxides, 
and by adding soluble carbonates to solutions of salts con- 
taining metals whose carbonates are insoluble. 

Examples. — The formation of potassium carbonate by 
the treatment of potassium hydroxide with carbon dioxide 
has already been illustrated. (See Experiment 94.) 

[Write the equation representing the action. Is the salt 
formed in this case soluble or insoluble in water?] 

The formation of calcium carbonate by passing carbon 
dioxide into a solution of calcium hydroxide (lime-water) 
has been illustrated under Carbon Dioxide. 

[Describe the experiment. Write the equation repre- 
senting the action in this case. Is calcium carbonate sol- 
uble or insoluble in water ? In hydrochloric acid, in sul- 
phuric acid, in nitric acid ? What action takes place with 
each of these acids ?] 



280 INTRODUCTION TO CHEMISTRY. 

Experiment 138. — The formation of carbonates by the addition 
of soluble carbonates to solutions of salts of metals whose carbon- 
ates are insoluble is illustrated by the following experiments: 
Make solutions of copper sulphate, iron sulphate, lead nitrate, 
silver nitrate, calcium chloride, barium chloride. Add to each a 
little of a solution of a soluble carbonate, as sodium carbonate, 
potassium carbonate, ammonium carbonate. Note the result in 
each case. Filter off all the precipitates, wash them thoroughly 
[Why ?J, and determine whether they are carbonates. This may 
be done by treating them with dilute acids, which decompose them, 
causing an evolution of carbon dioxide, which can be detected by- 
passing a little of it into lime-water. Write all the equations 
representing the reactions which take place in the above experi- 
ments. Here again, as in the experiments with the sulphates, the 
metals exchange places : 

C11SO4 + Na,CO a = Na 2 S0 4 + C11CO3. 

[Is copper bivalent or univalent if the formula of copper sul- 
phate is C11SO4?] 

All carbonates except those of the members of the potas- 
sium family are insoluble, and are decomposed by heat into 
carbon dioxide and the oxide of the metal. The decompo- 
sition of calcium carbonate into lime and carbon dioxide 
is the best-known illustration of this fact: 

CaC0 3 = CaO + C0 2 . 

Phosphates.— Calcium phosphate is very abundant in 
nature, and a few other phosphates are also found. The 
methods of making phosphates are in principle the same 
as those used in making sulphates. 

The phosphates of all the metals except the members of 
the potassium family are insoluble in water. The normal 
phosphates [What is a normal phosphate ?], as a rule, are 
not changed by heat. Those phosphates in which two 
thirds of the hydrogen is replaced by metal— as, for ex- 



PHOSPHA TES- SILICA TE8. 28 1 

ample, disodium phosphate, HNa 2 P0 4 — lose water when 
heated, and yield pyrophosphates: 

2HNa a P0 4 =* Na 4 P 9 7 + H 2 0. 

Sodium 
pyrophosphate. 

Those phosphates in which only one third of the hydro- 
gen is replaced by metal — as, for example, monosodium 
phosphate, H a NaP0 4 — lose water when heated, and yield 
metaphosphates : 

H,NaP0 4 = NaP0 3 + H,0. 

Sodium 
metaphosphate. 

Neither the pyrophosphates nor the metaphosphates are 
changed by heat. 

Silicates. — The extensive occurrence of silicates in nature 
has been spoken of. Those which are most abundant are 
the feldspars and their decomposition-products. The prin- 
cipal feldspar is a complex silicate of aluminium and potas- 
sium, of the formula KAlSi 3 O g , derived from the polysilicic 
acid H 4 Si 3 8 [What is a polysilicic acid?]: 

3H,Si0 3 = H,Si 3 6 + H,0. 

Silicates can be made by heating together at a high 
temperature silicon dioxide, in the form of fine sand, and 
bases. 

Experiment 139.— Mix together some fine sand and about four 
times its weight of a mixture of potassium and sodium carbon- 
ates. Heat in a platinum crucible in the flame of the blast-lamp* 
until the mass is thoroughly melted and the particles have com- 
pletely dissolved. Pour the molten mass out on a stone, and 
when cooled break it up and treat it with a little hot water. 
What passes into solution is a mixture of potassium and sodium 
silicates : 

Na 2 C0 3 + Si0 2 = Na 2 SiO s + C0 2 . 
* This is nothing but a large blowpipe worked by a foot-bellows.. 



282 INTRODUCTION TO CHEMISTRY. 

Some silicates are decomposed by the ordinary acids, 
such as sulphuric and nitric acids, the silicic acid separat- 
ing as a difficultly soluble substance, which loses water 
and becomes insoluble. 

Experiment 140. — Treat a little of the solution containing 
sodium and potassium silicates, prepared in the last experiment, 
with a little sulphuric or hydrochloric acid. A gelatinous sub- 
stance will be precipitated. This is silicic acid. Some of the acid 
remains in solution: 

Na 2 Si0 3 4- H 2 S0 4 = Na 2 S0 4 4- H 2 Si0 3 . 

By evaporating the solution to dryness and heating for a time 
on the water-bath, all the silicic acid is converted into silicon 
dioxide, which is entirely insoluble. 

Many silicates which are not acted upon by strong acids 
are decomposed when fused with sodium or potassium car- 
bonate. 

Silicates which are not decomposed in either of the ways 
mentioned yield to hydrofluoric acid. The action consists 
in the formation of the gas, silicon tetrafluoride, SiF 4 , 
and the fluorides of the metals present. Thus, the reaction 
in the case of feldspar takes place in accordance with the 
equation 

KAISi.O, + 16HF = KF + A1F 3 + 3SiF 4 + 8H 2 0. 

The silicon fluoride is given off and the fluorides of the 
metals are soluble in water. Hence hydrofluoric acid dis- 
solves the silicate. [Is this use of the word dissolves 
strictly correct ?] 



CHAPTER XX. 

THE POTASSIUM GROUP : LITHIUM, SODIUM, POTASSIUM, 
CESIUM, RUBIDIUM, (AMMONIUM). 

General. — The most-widely-distributed and hence best- 
known members of this group are sodium and potassium. 
The hypothetical metal ammonium is included in the 
group because the salts formed by ammonia, in which this 
hypothetical metal is considered to be present, very closely 
resemble the salts of potassium and sodium. The mem- 
bers of the group are generally called the metals of the al- 
kalies, as the two best-known members are obtained from 
the alkalies, caustic potash and caustic soda, or potassium 
and sodium hydroxides. 

Potassium, K (At. Wt. 39). — This element is a constitu- 
ent of many minerals, particularly of feldspar, which, as 
already explained, is a complex silicate of aluminium and 
potassium. It is found also in combination with chlorine 
as carnallite and sylvite; with sulphuric acid and alu- 
minium, as alum; with nitric acid, as saltpetre or potas- 
sium nitrate; and in other forms. The natural decomposi- 
tion of minerals containing potassium gives rise to the 
presence of this metal, in various forms of combination, 
everywhere in the soil. It is taken up by plants; and 
when vegetable material is burned the potassium remains 
behind, chiefly as potassium carbonate. When wood-ash 
is treated with water the potassium carbonate is dissolved, 
and it is obtained in an impure state by evaporating 

383 



284 INTRODUCTION TO CHEMISTRY. 

the solution. The substance thus obtained is called 
potash. 

Experiment 141. — Treat two or three litres of wood-ashes 
with water. Filter off the solution, and examine it by means of 
red litmus-paper. The color of the paper is changed to blue 
Plainly the solution is alkaline. Examine some potassium car- 
bonate. [Does its solution act in the same way ?J Evaporate 
the solution to dryness. Collect the dry residue and treat it 
with dilute hydrochloric acid. [Is a gas given off ? Is it carbon 
dioxide ?] 

Preparation. — The metal was first prepared by Sir Hum- 
phry Davy, in the year 1807, by the action of a powerful 
electric current on potassium hydroxide. It is now manu- 
factured by distilling a mixture of potassium carbonate 
and charcoal: 

K a C0 3 +2C = 2K + 3CO. 

Properties. — It is a light substance, which floats on 
water. [Have you had evidence of this ?] Its freshly-cut 
surface has a bright metallic lustre, almost white; it acts 
upon water with great energy, causing the evolution of hy- 
drogen, which burns, and the formation of potassium 
hydroxide. This reaction has already been considered in 
connection with hydrogen. [Turn back to the experiment 
(Experiment 27) and perform it again. It will now ap- 
pear much clearer.] In consequence of its action on 
water, potassium cannot be kept in the air. It is kept 
under some oil upon which it does not act, as petroleum. 

Compounds of Potassium. — The chief compounds of potas- 
sium with which we meet are the iodide, KI, which is ex- 
tensively used in medicine and in photography ; the hy- 
droxide, or caustic potash, KOH, which finds extensive use 
in laboratories; the nitrate, or saltpetre, KN0 3 , used in the 
manufacture of gunpowder; the chlorate, KC10 3 ,used iu 



POTASSIUM IODIDE-POTASSIUM HYDROXIDE. 285 

the preparation of oxygen and in medicine; and the car- 
bonate, K 2 C0 3 . 

The methods used in preparing some of these com- 
pounds are interesting, as illustrating the applications of 
the principles of chemistry. 

Potassium iodide, KI, is made by treating caustic potash 
with iodine until the solution begins to show a permanent 
yellow color, which is an indication that no more iodine 
will be taken up. The action is the same as that which 
takes place when chlorine acts upon warm concentrated 
caustic potash. Both the iodide and iodate are formed: 

6KOH + 61 = 5KI + GKIO3 + 3H,0. 

By evaporating off all the water and heating the residue, 
the iodate is decomposed into iodide and oxygen. 

Experiment 142. — Examine a bottle of crystallized potassium 
iodide. Taste a little. Dissolve some in water. Add some 
iodine to this solution. [Does the iodine dissolve ?] Heat a little. 
[Does it contain water of crystallization ?] Treat a crystal or two 
with a few drops of concentrated sulphuric acid. [What takes 
place ? To what is the appearance of violet vapors due ? How 
many gases are given off? (See Experiment 108.)] 

Potassium Hydroxide, KOH. — This well-known sub- 
stance, commonly called caustic potash, is prepared by 
treating potassium carbonate with calcium hydroxide in a 
silver or iron vessel. 

Experiment 143.— Dissolve 50 grams potassium carbonate in 
500 to 600 cc. water. Heat to boiling in an iron or silver vessel, 
and gradually add the slaked lime obtained from 25 to 30 
grams of good quick-lime. During the operation the mass should 
be stirred with an iron spatula. After the solution is cool, draw 
it off by means of a siphon into a bottle. This may be used in 
experiments in which caustic potash is required. 

The reaction is based upon the fact that calcium carbonate is 



986 INTRODUCTION TO CHEMISTRY. 

insoluble, and that potassium carbonate and calcium hydroxide 
are soluble : 

K 2 C0 3 -b Ca(OH) 2 = CaCOs + 2KOH. 

The hydroxide is a white brittle substance. In contact 
with the air it deliquesces [What does this mean ?] and ab- 
sorbs carbon dioxide. It is a very strong base. [Explain 
the action which takes place when potassium hydroxide 
acts upon ammonium chloride, NH 4 C1; copper sulphate, 
CuS0 4 ; and magnesium nitrate, Mg(N0 3 ) 2 . ] 

Potassium Nitrate, KN0 3 .— This salt is commonly called 
saltpetre. Its occurrence in nature has already been 
spoken of under Nitric Acid. [What are the conditions 
which give rise to its formation ?J When refuse animal 
matter is left to undergo decomposition in the presence of 
bases, nitrates are always the end-products. Advantage is 
taken of this fact for the purpose of preparing saltpetre 
artificially, the process being carried on on the large scale 
in the "saltpetre plantations." In these, refuse animal 
matter is mixed with earthy material, wood-ashes, etc., and 
piled up. These piles are moistened with the liquid prod- 
ucts from stables. After the action has continued for two 
or three years the outer crust is taken off and extracted 
with water. The solution thus obtained contains, besides 
potassium nitrate, calcium and magnesium nitrates. It is 
treated with a water-extract of wood-ashes or with potas- 
sium carbonate, by which the calcium and the magnesium 
are thrown down as carbonates. Much of the saltpetre 
which is in the market is made from Chili saltpetre, or 
sodium nitrate, by treating it with potassium chloride: 

NaN0 3 + KC1 ^ KN0 3 + NaCl. 

Potassium nitrate crystallizes in long rhombic prisms 
of salty taste. 

Uses of Potassium Nitrate. — It is used in the preparation 
of sulphuric acid [What part does it play in the prepara- 



GUNPOWDER. 287 

tion of sulphuric acid ?], and of nitric acid [How is nitric 
acid obtained from it ?]. Its chief use is in the manufac- 
ture of gunpowder. 

Gunpowder. — The value of gunpowder is due to the fact 
that it explodes readily, the explosion being a chemical 
change accompanied by a sudden evolution of gases. 
When the powder is enclosed in a gun-barrel the gases in 
escaping drive the ball before them. Gunpowder is made 
of a mixture of saltpetre, charcoal, and sulphur. When 
heated, the saltpetre gives off oxygen and nitrogen; the 
oxygen combines with the charcoal, forming carbon diox- 
ide and carbon monoxide, and the sulphur combines with 
the potassium, forming potassium sulphide. When a mix- 
ture of saltpetre and charcoal is burned, the reaction which 
takes place is this : 

2KNO, + 30 = CO, + CO + 2N + K 2 CO,. 

Experiment 144.— Mix together 15 grams potassium nitrate 
and 2.5 grams powdered charcoal. Set fire to the mass. 

[Problem. — What would be the volume, at 0° and under 760 
mm. pressure,of the gases evolved from 5 grams of gunpowder con- 
taining the constituents in exactly the proportions given above ?J 

By adding the necessary quantity of sulphur the carbon 
dioxide, which would otherwise remain in combination 
with the potassium as potassium carbonate, is given off and 
potassium sulphide formed: 

2KNO, + 3C + S = 3C0 2 + 2N + K 2 S. 

For this reaction the constituents should be mixed in 
the proportions: 

Saltpetre 74.83 

Charcoal 13.31 

Sulphur.: 11.86 



100.00 



288 INTRODUCTION TO CHEMISTRY. 

This is approximately the composition of all powder. 
When gunpowder explodes, the gases formed occupy about 
280 times the volume occupied by the powder itself. 

Potassium Chlorate, KC10 3 . — The reaction by which po- 
tassium chlorate is formed when chlorine acts upon a so- 
lution of potassium hydroxide has been discussed (see pp. 
109-113). In the manufacture of the chlorate it is found 
advantageous first to make calcium chlorate, and then to 
treat this with potassium chloride, when, at the proper con- 
centration, potassium chlorate crystallizes. The process in 
brief consists in passing chlorine into a solution of calcium 
hydroxide in which an excess of hydroxide is held in sus- 
pension. The first action leads to the formation of calcium 
hypochlorite. When the solution of this salt is boiled it is 
decomposed, yielding the chlorate and chloride: 

3Ca(OCl) 2 = Ca(0 3 01) 9 + 2CaCl 2 . 

On now treating the solution with potassium chloride 
the following reaction takes place : 

Ca(0 3 Cl) 2 + 2KC1 = 2KC10 3 + CaCl 2 . 

Properties. — Potassium chlorate gives up oxygen very 
easily and is hence a good oxidizing agent. It dissolves in 
water at the ordinary temperatures to the extent of 6 parts 
in 100 of water. 






Uses.— The chief uses of potassium chlorate are for the 
preparation of oxygen, and in the manufacture of matches 
and fireworks. The tips of Swedish safety-matches are 
made of potassium chlorate and antimony sulphide. The 
surface upon which they are rubbed to ignite them con- 
tains red phosphorus. The chlorate is extensively used in 
medicine, particularly as a gargle. 

Potassium Cyanide, KCN. — This salt is made by heating 
potassium ferrocyanide or yellow prussiate of potash, 



POTASSIUM SULPHATE— SODIUM. 289 

K 4 Fe(CN) 6 , with potassium carbonate, and extracting the 
mass with water. It is a violent poison. 

Potassium Sulphate, K 2 S0 4 . — This salt occurs in combi- 
nation with others in nature, particularly in the mineral 
kainite, which has the composition K 2 S0 4 .MgS0 4 .MgCl 2 -f- 
6H 2 0. This occurs in Stassfurt and in Kalusz. Potassium 
sulphate is used in medicine, and in the preparation of 
ordinary alum and of potassium carbonate. 

Sodium, Na (At. Wt. 23). — Sodium occurs very widely 
distributed and in large quantities, principally as sodium 
chloride. It is found in a number of silicates, and is a con- 
stituent of plants, especially of those which grow in the 
neighborhood of the sea-shore and in the sea. Just as the 
ashes of inland plants are rich in potassium carbonate, so 
the ashes of sea-plants and those which grow near the sea 
are rich in sodium carbonate. It is found everywhere in 
the soil, but generally in small quantity, its presence being 
due to the decomposition of minerals containing it, such as 
soda feldspar, or albite. It occurs also as sodium nitrate, 
and in large quantity in Greenland as cryolite, Na,AlF 6 , or 
AlF 3 .3XaF. 

Preparation. — It is prepared from sodium carbonate by 
the same method as that used in the preparation of potas- 
sium, the reaction involved being represented thus: 

Na,CO a + 2C = 2Na + 3CO. 

A method for the preparation of sodium has recently 
been devised by which its price has been materially lowered. 
This consists essentially in the reduction of sodium hy- 
droxide by heating it with an intimate mixture of finely- 
divided iron and carbon. The mass is prepared by mixing 
the iron with molten pitch, allowing it to cool, breaking it 
into pieces, and heating to a comparatively high tempera- 
ture without access of air. This is the Castner method. 



"290 INTRODUCTION TO CHEMISTRY. 

Properties. — Its properties are very similar to those of 
potassium. It is light, floating on water; it has a bright 
metallic lustre, and is soft. It decomposes water, but not 
as actively as potassium. 

[Describe what takes place when potassium is thrown 
upon water and when sodium is similarly treated. How is 
the difference accounted for?] 

Sodium readily unites with oxygen, and is used in some 
chemical processes as a reducing agent [What is a reducing 
agent?], as, for example, in the preparation of silicon, 
magnesium, and aluminium. A compound of mercury 
and sodium, known as sodium amalgam, is used in some 
metallurgical operations connected with the extraction of 
silver and gold from their ores. 

Compounds of Sodium. — The chief compounds of sodium 
are the chloride, NaCl; the hydroxide, or caustic soda, 
NaOH; the nitrate, or Chili saltpetre, NaNO a ; the sul- 
phate, Na 2 S0 4 ; the thiosulphate, Na 2 S 2 3 ; the carbonate, 
Na 2 C0 3 ; the borate, or borax, Na 2 B 4 7 ; the phosphate, 
HNa 2 P0 4 ; and the silicate, Na 2 Si0 3 . 

Sodium Chloride, NaCl. — This is the substance which is 
known as salt, or common salt. It occurs very widely dis- 
tributed, and, as it is easily soluble, much of the water 
which enters into the ocean contains some of it in solution. 
Sea-water contains from 2% to 3 per cent. The most im- 
portant deposits are those at Wieliczka in Galicia, at Stass- 
furt and Eeichenhall in Germany, and at Cheshire in 
England. Besides these there are, however, many other de- 
posits in the United States, in Africa, and in Asia. In 
some places the salt is taken out of mines in solid form; in 
others, water is allowed to flow into the mines, and to re- 
main for some time in contact with the salt, and the solu- 
tion thus formed is drawn or pumped out of the mine, and 
evaporated by appropriate methods. In hot countries salt 



SODIUM CHLORIDE-SODIUM NITRATE. 291 

is obtained by the evaporation of sea-water, the heat of the 
sun being used for the purpose. Large shallow cavities are 
made in the earth, and into these the water flows at high 
tide, or it is pumped up into them. 

Properties. — Sodium chloride crystallizes in colorless and 
transparent cubes. Sometimes that which occurs in nature 
is colored blue. In hot water it is but little more soluble 
than in cold water. In crystallizing, the crystals enclose 
water, not as water of crystallization, and this is given off 
when the crystals are heated, the action being accompanied 
by a crackling sound. This is known as decrepitation. 

Uses. — Salt is used as the starting-point in the prepara- 
tion of all sodium compounds and of chlorine and hydro- 
chloric acid. Salt is necessary to the life of man and many 
other animals. 

[How are chlorine and hydrochloric acid obtained from 
it? What takes place when a solution of silver nitrate is 
added to a solution of common salt ? What substances 
besides silver nitrate act in the same way ?] 

Sodium Hydroxide, NaOH. — This is commonly called 
caustic soda. It can be prepared in the same way as potas- 
sium hydroxide, that is, by treating a solution of sodium 
carbonate with lime. [Explain the reaction.] Its proper- 
ties are very similar to those of caustic potash. [Are the 
hydroxides of the metals mostly soluble or insoluble sub^ 
stances ?] 

Sodium Nitrate, NaN0 3 . — This compound occurs in 
large quantity in southern Peru, on the border of Chili, and 
is known as Chili saltpetre. The natural salt contains, 
besides the nitrate, sodium chloride, sulphate, and iodide. 
Sodium nitrate is very similar to potassium nitrate, but it 
cannot be used in place of the more expensive potassium 
salt in the manufacture of the finer grades of gunpowder, 



292 INTRODUCTION TO CHEMISTRY. 

as it becomes moist in the air and does not decompose 
quickly enough. It is used extensively in the manufacture 
of nitric acid, and also for the purpose of preparing ordi- 
nary saltpetre. The iodine contained in it is extracted on 
the large scale, and this forms an important source of 
iodine. 

Sodium Sulphate, Na 2 S0 4 + 10H 2 O. — The common name 
of this substance is Glauber's salt. It is manufactured in 
enormous quantities for the purpose of converting common 
salt into sodium carbonate, or " soda": 

2NaCl + HJS0 4 = Na 2 S0 4 + 2HC1. 

[What becomes of the hydrochloric acid which is given 
off?] 

The salt crystallizes in large colorless monoclinic prisms, 
containing 10 molecules of water of crystallization, Na 2 S0 4 
+ 10H 2 O. It loses water when left in contact with the 
air. [Is it" efflorescent or deliquescent ?] 

Sodium Thiosulphate, Na 2 S a O, + 5H 2 0.— This is the salt 
commonly called hyposulphite of soda. It is made on the 
large scale by treating caustic soda with sulphur, and con- 
ducting sulphur dioxide into the solution. It is also made 
by adding sulphur to a boiling solution of sodium sul- 
phite: 

Na 2 S0 3 + S = Na a S s O,. 

Its chief application is in photography, in which art it is 
used for the purpose of dissolving the excess of silver on 
the plate after exposure. 

Sodium Carbonate, Na 2 C0 3 + 10H 2 O. — This salt, com- 
monly called soda, is one of the most important of manu- 
factured chemical substances. The mere mention of the 
fact that it is essential to the manufacture of glass and soap 
will serve to give some conception of its importance. It is 



SODIUM CARBONATE-LE BLANC PROCESS. 293 

found in the ashes of sea-plants, just as potassium carbon- 
ate is found in the ashes of those plants which grow on the 
land. We are, however, not dependent on sea-plants for 
our supply, as two methods have been devised for prepar- 
ing sodium carbonate from sodium chloride with which 
nature provides us in such abundance. As these methods 
are interesting applications of chemical principles, it will 
be well to consider them briefly. 

The Le Blanc Process. — The problem to be solved is to 
convert sodium chloride, XaCl, into sodium carbonate, 
Na 2 C0 3 . The process devised by Le Blanc for the French 
government during the Revolution, when the supply had 
been cut off, involves four reactions: 

1st. The sodium chloride is converted into sodium sul- 
phate by treating it with sulphuric acid: 

2NaCl + H 2 S0 4 = Na a S0 4 + 2HC1. 

2d. The sodium sulphate thus obtained is heated with 
charcoal, which reduces it to sodium sulphide, Na 2 S: 

Na„S0 4 + 4C = Na a S + 4CO; 
Na a S0 4 + 2C = Na 2 S + 2C0 2 . 

3d. The sodium sulphide is heated with calcium car- 
bonate, when sodium carbonate and calcium sulphide are 
formed : 

Na 2 S + CaC0 3 = Na a CO a + CaS. 

Calcium sulphide is insoluble in water, so that by treat- 
ing the resulting mass w r ith water the sodium carbonate is 
separated from the sulphide. 

In practice the sodium sulphate is mixed with coal and 
calcium carbonate, and the mixture heated. The coal 
reduces the sulphate to the sulphide, which acts upon the. 



294 INTRODUCTION TO CHEMISTRY. 

calcium carbonate, forming sodium carbonate and calcium 
sulphide. The product of the action is known as crude 
soda or black ash. In order to purify this product, it is 
broken into pieces, and treated with water. Soda comes 
into the market as calcined purified soda, which contains 
no water of crystallization, and as crystallized soda, which 
has the composition Na 2 C0 3 -f 10H 2 O. 

The Solvay or Ammonia Process. — Another process which 
is extensively used is the so-called ammonia-soda process, 
or the Solvay process. This depends upon the fact that 
monosodium carbonate, HNaCO a , is comparatively diffi- 
cultly soluble in water. If, therefore, monoammonium car- 
bonate, or acid ammonium carbonate, HNH 4 C0 3 , is added to 
a solution of common salt, acid sodium carbonate, HNaC0 3 , 
crystallizes out, and ammonium chloride remains behind in 
the solution: 

NaCl + HNH 4 C0 3 = HNaC0 3 + NH 4 C1. 

When the acid carbonate is heated, it gives off carbon 
dioxide, and is converted into the normal salt thus: 

2HNaC0 3 = tta 2 C0 3 +H 2 0. 

The carbon dioxide given off is passed into ammonia, 
and thus again obtained in the form of acid ammonium 
carbonate : 

NH, + H 2 + C0 2 - HNH 4 C0 3 . 

The ammonium chloride obtained in the first reaction is 
treated with lime or magnesia, MgO, and the ammonia set 
free. This ammonia is again used in the preparation of 
acid ammonium carbonate. 

More than half the soda supply of the world is now fur- 
nished by the Solvay process, 



DISODIUM PHOSPHATE. 295 

Experiment 145. — Make a solution of common salt in ordinary 
ammonia-water (about 50 cc). Pass carbon dioxide into this 
solution. Use a funnel attached to the delivery tube as in Exp. 
59. Filter off the precipitate formed, and dry it by spreading it 
upon layers of filter-paper. Heat some of the salt when dry, and 
determine whether the gas given off is carbon dioxide or not. 
When gas is no longer given off by heat, let the tube cool and 
examine the residue. [Is it a carbonate ?] 

Properties. — Sodium carbonate crystallizes in large mono- 
clinic prisms with 10 molecules of water of crystallization. 
The crystals are efflorescent. 

Monosodium Carbonate, Primary Sodium Carbonate, 
HNaCO s . — This salt is commonly called " bicarbonate of 
soda." It is easily prepared by passing carbon dioxide 
over the ordinary carbonate dissolved in its water of crys- 
tallization : 

Na a CO, + C0 2 +H a O .= SHNaCO,. 

When heated it gives up carbon dioxide and water, and 
is converted into the normal salt. It is used in medicine, 
and extensively in the preparation of soda-water and other 
effervescent drinks. 

Disodium Phosphate, HNa a P0 4 + 12H a O.— This is the 
common form of sodium phosphate. It is formed when 
phosphoric acid is treated with sodium carbonate until the 
solution begins to show an alkaline reaction with red lit- 
mus. It is a remarkable fact that, although phosphoric 
acid is tribasic, and with most metals forms salts that are 
derived from the acid by replacement of all the three 
hydrogen atoms, as Ag 3 P0 4 , Ca 3 (P0 4 ) 2 , etc., with sodium 
its most stable salt is the one in which two hydrogen atoms 
are replaced by sodium. A salt of the formula Na 3 P0 4 
can be made, but it has an alkaline reaction, and absorbs 



296 INTRODUCTION TO CHEMISTRY. 

carbon dioxide from the air, being converted into sodium 
carbonate and disodium phosphate; 

21\a 3 P0 4 +■ C0 2 + H 2 = 2HNa 2 P0 4 + Na 2 C0 3 . 

Sodium Borate, Na 2 B 4 7 + 10H 2 O.— This salt has been 
referred to under Boric Acid. It is commonly called borax. 
It is found in nature in several lakes in Asia, and in this 
country in Clear Lake, Nevada. It is manufactured by 
neutralizing the boric acid found in Tuscany. 

When heated, borax puffs up, and at red heat melts, 
forming a transparent colorless liquid. This is anhydrous 
borax, Na 2 B 4 7 . Molten borax has the power to dissolve 
metallic oxides, and forms colored glasses with some of 
them. It is used in blowpipe work (see Boric Acid). As 
it dissolves metallic oxides, it is used in the process of 
soldering, as it is necessary to have bright, untarnished 
metallic surfaces in order that the solder shall adhere 
firmly. Borax is an antiseptic; that is to say, it prevents 
the decomposition of organic substances. It is used exten- 
sively in the manufacture of porcelain and in glass-paint- 
ing. 

Ammonium Salts. — The method of formation of the 
so-called ammonium salts has been described (see Ammo- 
nia). These salts resemble the salts of potassium and 
sodium in many respects, and they are hence described in 
the same connection. The chief ones are the chloride, 
NH 4 C1; the carbonate, (NH 4 ) 2 C0 3 ; the sulphide, (NH 4 ) 2 S, 
the hydrosulphide, (NH 4 )HS; and sodium-ammonium phos- 
phate, HNaNH 4 P0 4 + 4H 2 0. 

Ammonium Chloride, NH 4 C1. — This salt is commonly 
called sal ammoniac. At present its principal source is 
the gas-works. The ammonia-water of the works is neu- 
tralized with hydrochloric acid, and the salt obtained by 
evaporation. It has a sharp, salt taste, and is easily soluble 



AMMONIUM SULPHIDE. 297 

in water. When heated it is converted into vapor without 
melting, and with very slight decomposition; and when 
the vapor comes in contact with a cold surface, it con- 
denses in the form of crystals. This process of vaporizing 
and condensing a solid is called sublimation. 

Experiment 146.— On a piece of platinum foil or porcelain heat 
a little pure ammonium chloride. It will pass off and form a 
dense white cloud. This is the same cloud as that formed by 
bringing together gaseous ammonia and hydrochloric acid. All 
ammonium salts are either volatile or decompose when heated. 

[What takes place when ammonium chloride is treated 
with caustic soda ? with lime ? with sulphuric acid ?] 

Ammonium Sulphide, (NH 4 ) 2 S. — This substance is exten- 
sively used in chemical analysis for the purpose of precipi- 
tating those sulphides which are soluble in dilute hydro- 
chloric acid. As will be remembered, in analyzing a 
mixture of substances the first thing usually done is to add 
hydrochloric acid to the solution. This precipitates silver, 
lead, and, under certain conditions, mercury. This pre- 
cipitate having been filtered off, hydrogen sulphide is 
passed through the filtrate, when those metals whose sul- 
phides are insoluble in dilute hydrochloric acid are thrown 
down. The precipitate is filtered off and ammonium sul- 
phide added to the filtrate, when the metals whose sul- 
phides are soluble in dilute hydrochloric acid are thrown 
down. Among these are iron, cobalt, nickel, manganese, 
etc. Any other soluble sulphide might be used, but the 
advantage of ammonium sulphide is that it is volatile, and, 
hence, by evaporating the solution and heating, it can be 
got rid of. 

Ammonium sulphide is made by passing hydrogen sul- 
phide into an aqueous solution of ammonia. If the gas is 
passed until the solution is saturated, the product is the 
hydrosulphideHNHJS: 

NH, + H a S = HNHJS. 



298 INTRODUCTION TO CHEMISTRY. 

If only half this quantity of the gas is passed, the 
product is the sulphide : 

2NH 3 + H 2 S = (NH 4 ) 2 S. 

The simplest way to make it is to divide a quantity of 
ammonia solution into two equal parts. Saturate one half, 
thus forming the hydrosulphide, and add the other half, 
when this reaction takes place : 

HNH 4 S + NH 3 = (NH 4 ) 2 S. 

The product is a colorless liquid of a disagreeable odor. 
It soon changes color, becoming yellow, and after a time a 
yellow deposit is formed in the vessel in which it is con- 
tained. This change of color is due to the action of the 
oxygen of the air. Some of the sulphide is decomposed 
into ammonia, water, and sulphur: 

(NH 4 ),S + = 2NH, + H a O + S. 

The sulphur thus set free combines with the unde- 
composed ammonium sulphide, forming the compounds 
(NH 4 ) 2 S 2 , (NH 4 ) 2 S 3 , etc., known as poly sulphides. When 
as much sulphur as possible has been taken up in this way, 
any more which may be set free by the action of oxygen is 
deposited, 

A solution containing the polysulphides is called yellow 
ammonium sulphide. It is used to dissolve the sulphides 
of arsenic, antimony, and tin in analytical operations. (See 
description of method of analysis, p. 418.) 

Experiment 147. — Saturate 50 cc. strong aqueous ammonia 
. with hydrogen sulphide. Add to the saturated solution 50 cc. of 
the same ammonia. 

Ammonium Hydrosulphide, HNH 4 S. — As stated above, 
a solution of this substance is made by passing hydrogen 



CHARACTERISTICS OF METALS OF THE ALKALIES. 299 

sulphide into a solution of ammonia until no more is 
taken up. 

Sodium-ammonium Phosphate, HNaNH 4 P0 4 + 4H 2 0. — 
This is commonly called microcosmic salt, and is much 
used in the laboratory in blowpipe work. Its value in this 
kind of work depends upon the fact that it is decomposed 
by heat, yielding sodium metaphosphate: 

HNaNH 4 P0 4 = NaP0 3 + NH 3 + H 2 0; 

and the metaphosphate at high temperatures combines 
with the metallic oxides, forming double phosphates, many 
of which are colored. 

General Characteristics of the Metals of the Alkalies. — 
From what has been said, it will be seen that nearly all 
the compounds of these metals are soluble in water. Of 
those mentioned only monosodium carbonate is at all 
difficultly soluble. There are a few insoluble salts of 
potassium, those which are chiefly used in analytical opera- 
tions being the chloroplatinate, K 2 PtCl 6 , which is formed 
by adding a solution of platinum chloride, PtCl 4 , to a 
solution containing potassium chloride: 

2KC1 + PtCl 4 = K 2 PtCl 6 ; 

and the fluosilicate, K 2 SiF 6 , which is formed when a solu- 
tion of fluosilicic acid, H 2 SiF 6 , is added to a solution con- 
taining a salt of potassium. 

Experiment 148. — Add platinum chloride and fluosilicic acid 
successively to solutions containing potassium chloride. If in 
the former case no precipitate is formed add a little alcohol ; 
potassium chloroplatinate is slightly soluble in water, but is insol- 
uble in dilute alcohol. Potassium fluosilicate is precipitated only 
from rather concentrated solutions, and even from these, as a 
rule, only after standing for a while. 



300 INTRODUCTION TO CHEMISTRY. 

Rare Elements of this Group.— The elements lithium, 
ccesium, and rubidium are much rarer than sodium and 
potassium. Lithium is found in a form of mica known as 
lepidolite. It is the lightest metal known, and has the 
smallest atomic weight, viz., 7. 

Relations between the Atomic Weights of the Members 
of this Group. — The relations between the atomic weights 
of the members of this group are similar to those already 
noticed between chlorine, bromine, and iodine; sulphur, 
selenium, and tellurium; and phosphorus, arsenic, and 
antimony. Thus, we have lithium, 7 ; sodium, 23; and 
potassium, 39. The atomic weight of sodium, 23, is the 
mean of those of lithium, 7, and potassium, 39: 

2 - M " 

Similarly, the atomic weight of rubidium, 85, is nearly 
the mean of those of potassium and caesium, 133: 

39 + 133 

5 = 86 - 

Flame Reactions. — When a clean piece of platinum wire 
is held for some time in the flame of the Bunsen burner, 
it then imparts no color to the flame. If now a small piece 
of sodium carbonate or any other salt of sodium is put on 
it, the flame is colored intensely yellow. All sodium com- 
pounds have this power, and the chemist makes use of the 
fact for the purpose of detecting the presence of sodium. 
Similarly, potassium compounds color the flame violet ; 
lithium compounds color the flame red; and the other 
metals of the family also impart characteristic colors to 
the flame. 

Experiment 149. — Prepare some pieces of platinum wire, 8 to 
10 cm. long, with a loop on the end, like those described for 



THE SPECTROSCOPE. 301 

blowpipe work. After thoroughly cleaning them, insert one in a 
little sodium carbonate, and notice the color it gives to the flame. 
Try another with potassium carbonate, and, if the substances are 
available, others with a lithium, a caesium, and a rubidium com- 
pound. 

The Spectroscope. — While it is an easy matter to rec- 
ognize potassium alone, or any one of the other metals 
alone, it is difficult to do so when they are together in the 
same compound. For example, when potassium and 
sodium are together, the intense yellow caused by the 
sodium completely masks the more delicate violet caused 
by the potassium, so that the latter cannot be seen with 
the unaided eye. In this particular case we can get over 
the difficulty by letting the light pass through a blue 
glass, or a thin glass vessel filled with a solution of indigo. 
The yellow light is thus cut off, while the violet light 
passes through and can be recognized. A more general 
method for detecting the constituents of light is by means 
of a prism. Lights of different colors are turned out of 
their course to different extents when passed through a 
prism, as is seen when white sunlight is passed through a 
prism. A narrow beam of white light passing in emerges 
as a band of various colors, called its spectrum. We thus 
see that white light is made of different colored lights. 
Similarly, we can determine what any light is composed of. 
Every light has its own characteristic spectrum. The 
light produced by burning sodium, or by introducing a 
sodium compound in a colorless flame, has a spectrum 
consisting of a narrow yellow band. The spectrum of potas- 
sium consists essentially of two bands, one red and one 
violet. Further, these bands always occupy definite posi- 
tions relatively to one another, so that, on looking through 
a prism at the light caused by potassium and sodium, the 
yellow band of sodium is seen in its position, and the two 
potassium bands in their proper positions. 

The instrument used for the purpose of observing the 



302 INTRODUCTION TO CHEMISTRY. 

spectra of different lights is called the spectroscope. It 
consists essentially of a -prism and two tubes. Through 
one of the tubes the light to be examined is allowed to 
pass so as to strike the prism properly. The light emerges 
from the other side of the prism, and is observed through 
the other tube, which is provided with lenses for the 
purpose of magnifying the spectrum. By means of the 
spectroscope it is possible to detect the minutest quantities 
of some elements, and, since it was devised, several new 
elements have been discovered through its aid, as, for 
example, caesium, rubidium, thallium, indium, gallium, and 
others.* 

* For an account of the spectroscope and its uses tbe student is 
advised to consult some work on physics. The principles involved 
in its construction are physical principles, and cannot properly be 
taken up in detail in a text-book of chemistry. 



CHAPTER XXI. 

THE CALCIUM GROUP: 
CALCIUM, BARIUM, STRONTIUM, GLUCINUM. 

General. — The three elements calcium, barium, and 
strontium resemble one another very closely. Calcium is 
much more abundant than either of the other members of 
the group, while strontium is the least abundant of the 
three. For the present it will be best to confine our atten- 
tion to the principal member, viz., calcium. 

Calcium, Ca (At. Wt. 40). — This element occurs very 
widely distributed in nature, and in enormous quantities. 
It is found principally as carbonate, CaCO,, in the form of 
limestone, marble, and chalk; as sulphate, CaS0 4 , in the 
form of gypsum; as phosphate, Ca 3 (P0 4 ) a , in phosphorite 
and apatite; as fluoride, CaF 2 , in fluor-spar. 

The element is made by decomposing calcium chloride 
by means of the electric current. 

It is a brass-yellow lustrous substance, which in moist 
air becomes covered with a layer of hydroxide. At ordi- 
nary temperatures it decomposes water just as sodium and 
potassium do. 

Compounds of Calcium. — The principal compounds of cal- 
cium with which Ave have to deal are the chloride. CaCl 3 : 
the oxide, or quick-lime, CaO; the hydroxide, or slaked 
lime, Ca(OH) 2 ; the hypochlorite, Ca(OCl) 2 ; the carbonate, 
CaC0 3 ; the sulphate, CaS0 4 ; the jjliosphate, Ca 3 (P0 4 ) 2 ; 
and the silicate, in the form of glass. 

303 



304 INTKODUCTION TO CHEMIST BY. 

Calcium Chloride, CaCl 2 . — The property which gives this 
salt its value is its power to absorb water. It is used as a 
drying agent. Gases are passed through it for the purpose 
of drying them, and it is also placed in vessels in which it 
is necessary that the atmosphere should be dry. 

Experiment 150. — Dissolve 10 to 20 grams of limestone or 
marble in ordinary hydrochloric acid. Evaporate to dryness. 
Expose a few pieces of the residue to the air. Does it become 
moist? In what experiments has calcium chloride been used, 
and for what purposes ? What would happen if sulphuric acid 
were added to calcium chloride? Try it. Explain what takes 
place. Is the residue soluble or insoluble in water ? 

Calcium Oxide, CaO. — This is the substance commonly 
called lime, or, to distinguish it from the hydroxide or 
slaked lime, it is called quick-lime. It is made by heating 
calcium carbonate, which is thus decomposed into lime and 
carbon dioxide : 

CaC0 3 = CaO+C0 2 . 

[In what connection have we already met with this re- 
action ?] 

Limekilns are large furnaces in which limestone and 
other forms of calcium carbonate are heated and converted 
into lime. 

[Why is it dangerous to remain for any length of time in 
the immediate neighborhood of a limekiln ?] 

Lime is a white, amorphous, infusible substance. When 
heated in the flame of the compound blowpipe, it gives 
forth an intense light, as any other infusible substance 
would under the same circumstances. When exposed 
to the air, it attracts moisture and carbon dioxide and 
is thus converted into the carbonate. It must hence 
be protected from the air. Lime which has been con- 
verted into the carbonate by exposure to the air is said to 
be air-slaked, 



CALCIUM HYDROXIDE. 305 

Calcium Hydroxide, Ca0 2 H 2 , or Ca(0H) a — When cal- 
cium oxide or quick-lime is treated with water, it be- 
comes hot and crumbles to a fine powder. The substance 
formed in this operation is somewhat soluble in water, 
the solution being known as lime-water. The chemical 
change that takes place when lime is treated with water 
has been explained. It consists in the formation of a 
compound of the formula Ca0 2 H 2 and known as slaked 
lime, and the operation is known as slaking. It is be- 
lieved that just as potassium hydroxide, KOH, is properly 
regarded as water in the molecule of which one atom of 
hydrogen is replaced by an atom of potassium, so calcium 
hydroxide is properly regarded as derived from water by 
the replacement of two atoms of hydrogen in two mole- 
cules by one atom of the bivalent metal calcium: 

HOH r \ OH n /nm 

HOH Ca \ OH or Ca (OH) 9 . 

Two mol. water. Calcium hydroxide. 

It is difficult to explain exactly why this view is held. 
It can only be said that it is a conception in harmony with 
a great many facts, though it does not follow as a nec- 
essary consequence from any facts known to us. 

Experiment 151. — To 40 to 50 grams good quick-lime add 
100 cc. water. Soon the mass will begin to crumble, and steam 
will rise from it, indicating that heat is evolved. Afterwards 
dilute to 2 to 3 litres and put the whole in a well-stoppered bot- 
tle. The undissolved lime will settle to the bottom, and in the 
course of some hours the solution above will become clear. 
Carefully pour off some of the clear solution. [What takes place 
when some of the solution is exposed to the air? when the 
gases from the lungs are passed through it ? when carbon di- 
oxide is passed through it ? What takes place when dilute 
sulphuric acid is added to lime-water ? Is calcium sulphate 
difficultly or easily soluble in water? Has lime-water an alka- 
line reaction ?] 



306 INTRODUCTION TO CHEMISTRY. 

When potassium hydroxide is added to a solution of a 
salt containing a metal whose hydroxide is insoluble in 
water, the insoluble hydroxide is precipitated. This was 
illustrated in Experiments 130 and 131. [What are those 
experiments ?] Calcium hydroxide is a soluble hydroxide, 
and acts in the same way that potassium hydroxide does. 

Experiment 152. — Add some lime-water to a dilute solution of 
ferric chloride, of copper nitrate, of lead nitrate. Explain the 
results. 

Uses. — Lime is extensively used in the arts, generally in 
the form of the hydroxide. As we have seen, it is used in 
the preparation of ammonia and of the caustic alkalies, po- 
tassium and sodium hydroxides; and of bleaching-powder 
and potassium chlorate. It is further used in large quan- 
tity in the process of tanning for removing the hair from 
hides; in decomposing fats for making stearin for candles; 
for purifying illuminating-gas; and especially in the prep- 
aration of mortar. 

Calcium hypochlorite, Ca(OCl) 2 , has already been refer- 
red to under the head of Chlorine. The form in which 
chlorine is transported is " bleaching-powder," a com- 
pound containing calcium hypochlorite and calcium chlo- 
ride, Ca(OCl) 2 + CaCl 2 , made by treating slaked lime with 
chlorine : 

2Ca(OH) 2 + 401 = Ca(OCl) 2 + CaCl 2 + 2H a O. 

Bleaching-powder. 

The compound is commonly called " chloride of lime." 
An objection to the view that calcium chloride is present 
as such in bleaching-powder is found in the fact that the 
substance is not deliquescent. This has led to the sug- 
gestion that bleaching-powder in the dry form is not a 
mixture of two compounds as represented above, but that 



BLEA CHING PO WDER. 3( Yl 

CI 
it is rather one compound of the formula Ca<^p, or 

CaOCl 2 . The point is a difficult one to decide, but at 
present the evidence appears to be in favor of the view 
that bleaching-powder in the dry form is a single com- 
pound of the constitution represented by the formula 
last given. When treated with water, however, it appears 
to be resolved into a mixture of the hypochlorite and 
chloride. 

Properties. — Bleaching-powder is a white substance that 
has the odor of hypochlorous acid. When treated with an 
acid it gives up all its chlorine. When exposed to the 
action of carbon dioxide hypochlorous acid is liberated. 
Hence this decomposition takes place slowly in the air. 

How Bleaching-powder Acts in Bleaching. — A solution of 
bleaching-powder alone is not capable of bleaching except 
very slowly. If, however, something is added which has 
tiie power to decompose it, bleaching takes place, the 
action being due to the presence of hypochlorous acid and 
chlorine. As is clear from what was said above, the 
passage of carbon dioxide through the solution or the 
addition of an acid would cause it to bleach. So, too, 
certain salts produce a similar effect. The explanation of 
this is the instability of the hypochlorites formed by the 
skits added. 

Decomposition of Bleaching-powder by Boiling its Solu- 
tion. — When a concentrated solution of bleaching-powder 
is heated it gives off oxygen, and the salt is converted into 
the chloride. In dilute solution, however, the hypochlorite 
is converted into chlorate and chloride: 

3Ca(C10), = Ca(C10 3 ), + 2CaCl 2 . 



308 INTRODUCTION TO CHEMISTRY. 

This fact is taken advantage of, as has been shown, for 
the purpose of making calcium chlorate, and from this 
potassium chlorate (see p. 288). In contact with certain 
oxides, as copper oxide, ferric oxide, and with hydroxides, 
as those of cobalt and nickel, a solution of bleaching- 
powder readily gives up oxygen when heated. 

Uses. — The chief application of bleaching-powder is, as 
its name implies, for bleaching. It is also used as a disin- 
fectant and as an antiseptic, that is, for the purpose of 
destroying disease germs and of preventing decomposition 
of organic substances. 

Calcium Carbonate, CaC0 3 . — This salt occurs in im- 
mense quantities in nature in the well-known forms lime- 
stone, calc-spar, marble, and chalk. The variety of calc- 
spar found in Iceland, and known as Iceland spar, is par- 
ticularly pure calcium carbonate. It crystallizes in a 
number of different forms, the most common being in 
rhombohedrons, as seen in ordinary calc-spar. A second 
variety of crystallized calcium carbonate is aragonite. 
This is found in nature crystallized in rhombic prisms, 
and in forms derived from this. AVhen heated, aragonite 
falls to pieces, the particles being small crystals of the form 
characteristic of calc-spar. This is a case of dimorphism 
similar to that presented by sulphur, which, it will be 
remembered, crystallizes in two forms, rhombic and mono- 
clinic, the latter of which passes into the former spon- 
taneously. These forms are produced artificially very 
readily. AVhen calcium carbonate is precipitated from a 
solution of a calcium salt by adding a soluble carbonate at 
ordinary temperatures, the precipitate is made up of micro- 
scopic crystals which have the same form as calc-spar. If, 
however, the solution from which the carbonate is precipi- 
tated is hot, the salt consists of microscopic crystals of the 
form of aragonite. 



CALCIUM CARBONATE. 3<)9 

The most abundant form of calcium carbonate is lime- 
stone, of which many great mountain-ranges are largely 
made up. This is a compact form of the compound, which 
has a gray color, and frequently consists of minute crystals. 
It is always more or less impure, containing clay and other 
substances. Limestone which is mixed with a considerable 
proportion of clay is called marl. Many natural waters 
contain calcium carbonate in solution — probably in the 
form of the acid carbonate. When such a water evapo- 
rates, the carbonate is deposited. It happens in some 
places that a water charged with the carbonate works its 
way slowly through the earth and drops from the top of a 
cave. Under these circumstances there is a gradual 
deposit of the salt which remains suspended. Such hang- 
ing formations of the carbonate are known as stalactites. 
At the same time that part of the liquid which falls to the 
bottom of the cave forms a projecting mass below the 
stalactite. Such projecting masses are called stalagmites. 
The formation of stalactites takes place in much the same 
way as that of icicles. 

Much of the calcium carbonate found in nature has its 
origin in the remains of animals, and fossils are very 
abundant in it. Chalk consists almost exclusively of the 
shells of microscopic animals. 

Temporary Hardness. — When carbon dioxide is passed 
into a solution of calcium hydroxide, the carbonate is pre- 
cipitated; and, if the current of gas is continued long 
enough, the carbonate is redissolved. On heating the so- 
lution to boiling, the normal carbonate is precipitated and 
carbon dioxide is given off. Natural w r aters w r hich come in 
contact with limestone gradually take up more or less of 
the carbonate, with the aid of the carbon dioxide of the air, 
and when such a water is boiled, the carbonate is thiwvn 
down. A w r ater containing calcium carbonate in solution is 
called a hard water; and, as this kind of hardness is easily 



310 INTRODUCTION TO CHEMISTRY. 

removed by boiling, it is called temporary hardness in order 
to distinguish it from a kind which is not removed by boil- 
ing, and is therefore called permanent hardness. Further, 
temporary hardness is removed by adding lime to the 
water, when the normal carbonate is formed, which is at 
once precipitated. 

The decomposition of calcium carbonate by heat, form- 
ing lime, or calcium oxide, and carbon dioxide, was referred 
to on p. 304. 

Applications. — Calcium carbonate is used, in the arts, 
for a great many purposes, as in the manufacture of glass; 
as a flux in many important metallurgical operations, 
as in the reduction of iron from its ores; in the prep- 
aration of lime for mortar, etc. As is well known, 
further, marble and some of the varieties of limestone are 
extensively used in building; and large quantities of chalk 
are also used. 

Calcium Sulphate, CaS0 4 .— This compound is very 
abundant in nature. The principal natural variety is 
gypsum, which occurs in crystals containing two molecules 
of water, CaS0 4 + 2H 2 0. The salt of the formula CaS0 4 
also occurs in nature, and is called anhydrite. A granular 
form of gypsum is called alabaster. Calcium sulphate is 
difficultly soluble in hot and cold water. When heated to 
100°, or a little above, it loses nearly all of its water and 
forms a powder known as plaster of Paris, which has the 
power of taking up water and forming a solid substance. 
This process of solidification is known as "setting." 
Plaster of Paris is very largely used in making casts, on 
account of its power to harden after having been made 
into a paste with water. The hardening is a chemical 
process, and is caused by the combination of water with 
the salt to form the crystallized variety. 

When heated to 200°, and above, all the water is given 



SOLUBLE CARBONATES AND GYPSUM 311 

off from gypsum, and the product now combines with 
water only very slowly, and is of no value for making casts. 
In general, the higher the temperature to which the gyp- 
sum is heated, the greater the difficulty with which the 
product combines with water. 

Experiment 153. — Heat some powdered gypsum to about 200° 
in an air-bath. Examine the residue and see whether it will 
jcc-ome solid when mixed with a little water so as to form a paste. 

Permanent Hardness. — Many natural waters contain 
gypsum in solution. Such waters act in some respects like 
those which contain calcium carbonate. With soap, for 
example, they form insoluble compounds. This kind of 
hardness is not removed by boiling, and it is therefore 
called permanent hardness. Magnesium sutyhate acts in 
the same way, producing permanent hardness. 

Action of Soluble Carbonates on Gypsum. — When calcium 
sulphate is treated with a solution of a soluble carbonate, 
it is decomposed, forming calcium carbonate as represented 
in the equation 

CaS0 4 + Na 2 C0 3 = Na 2 S0 4 + CaC0 3 . 

This change is effected simply by allowing the two to 
stand in contact at the ordinary temperature. 

Experiment 154. — Upon a gram or two of powdered gypsum 
pour, say, 50 cc. of a moderately strong solution of ammonium 
carbonate. After a few hours pour off the solution, collect the 
powder on a filter, wash it thoroughly with water [Why ?J and 
see whether it has changed to the carbonate. [How can you de- 
termine whether ammonium sulphate is in solution or not ? Of 
course, there is still ammonium carbonate present, and this must 
be taken into account in examining for the sulphate. We usually 
examine for a sulphate by adding a soluble barium salt, when, if 
a soluble sulphate is present, barium sulphate is precipitated. In 
this case, however, the ammonium carbonate w 7 ould throw down 



31 2 INTRODUCTION TO CHEMISTRY. 

barium carbonate. To prevent this, the ammonium carbonate is 
decomposed by slowly adding sufficient dilute hydrochloric acid. 
There will then be present ammonium chloride and sulphate ; 
and, now, if barium chloride or any other soluble barium salt is 
added, barium sulphate is precipitated.] 

Uses. — Besides its rise for making casts, calcined gypsum 
is used also in surgery for making plaster-of-Paris ban- 
dages, and as a fertilizer. Its action as a fertilizer is 
thought by some to be due to the fact that it has the power 
to hold ammonia and ammonium carbonate in combina- 
tion, and thus to make them available for the plants. It 
has been shown that it in some way facilitates the process 
of nitrification, and perhaps it is in consequence of this 
that it aids plant-growth. 

Calcium Phosphates. — There are three phosphates of 
calcium: (1) the normal phosphate, Ca :) (P0 4 ) 2 ; (2) the 
secondary phosphate, CaHP0 4 ; and (3) the primary phos- 
phate, CaH 4 (P0 4 ) a . 

Normal calcium phosphate is found in nature in large 
quantity as phosphorite, and in combination with calcium 
fluoride or chloride as apatite. It is, further, the principal 
inorganic constituent of bones, forming 85 per cent of 
bone-ash. 

Calcium Phosphate Essential to Plant-growth. — It is 

found everywhere in the soil, and is taken up by the plants 
for whose growth it is essential. That it is also essential 
to the life of animals is obvious from the fact that the 
bones consist so largely of it. The phosphate required for 
the building up of bones is taken into the system with the 
food. From these statements it is clear that calcium phos- 
phate is of fundamental importance, and that a fertile soil 
must either contain this salt, or something from which it 
can be formed. Now, when a crop is raised on a given 
area, a certain amount of the phosphate contained in it is 



ARTIFICIAL FERTILIZERS. 313 

withdrawn. If the plants were allowed to decay where 
they grow, the phosphate would be returned and the soil 
would continue fertile; but in cultivated land this is not 
the case. The crops are removed, and with them the cal- 
cium phosphate, and the soil therefore becomes exhausted. 
If the substances removed are used as food, some of the 
phosphate is found in the excrement of the animals; and, 
if the excrement is put on .the soil, this is again rendered 
fertile. 

Artificial Fertilizers. — There are, however, other sources 
of calcium phosphate, and some of these are utilized exten- 
sively in the preparation of artificial fertilizers. The nat- 
ural form of the phosphate, as that in bone-ash, in phos- 
phorite, and in guano, is mainly the normal or neutral 
phosphate. This is insoluble in wrter, and is therefore 
taken . up by the plants with difficulty. To make it 
quickly available, it must be converted into a soluble phos- 
phate. This is done by treating it with sulphuric acid in 
order to effect the reaction represented in this equation : 

Ca 3 (P0 4 ) 2 + 2H 2 S0 4 = CaH 4 (P0 4 ) 2 + 2CaS0 4 . 

The primary phosphate thus formed is soluble in water, 
and is of great value as a fertilizer. The mixture of the 
soluble phosphate and of calcium sulphate is known as 
" superphosphate of lime." The sulphate, as we have seen, 
is also of value as a fertilizer. The value of superphos- 
phates depends mostly upon the amount of soluble phos- 
phate contained in them; and in dealing with them it is 
customary to state how much " soluble " and how much 
" insoluble phosphoric acid " they contain. When a super- 
phosphate is allowed to stand for a time, some of the 
soluble primary phosphate is converted into insoluble 
phosphates by contact with basic hydroxides and water. 
This is known as the process of " reversion," and that part 



314 INTRODUCTION TO CHEMISTRY. 

of the phosphoric acid which is contained in the insoluble 
phosphate is spoken of as " reverted phosphoric acid." 

Formation of Calcium Phosphate by Precipitation. — 

Normal calcium phosphate, as has been stated, is insoluble 
in water, and is formed when a soluble normal phosphate 
is added to a solution of a calcium salt. It is also formed 
when disodium phosphate and ammonia are added to a 
solution of a calcium salt, thus: 

2HNa 2 P0 4 + 3CaCl a + 2NH 3 = Ca 3 (P0 4 ) 2 + 4NaCl + 2NH 4 C1. 

Experiment 155. — To a solution of calcium chloride in a test- 
tube add disodium phosphate and ammonia. The precipitate will 
dissolve in hydrochloric or nitric acid. 

Primary calcium phosphate, CaH 4 (P0 4 ) 2 , is commonly 
called the acid phosphate of calcium. It is formed when 
ordinary insoluble calcium phosphate is treated with con- 
centrated sulphuric acid, and is contained in the so-called 
superphosphates. 

Calcium silicate, CaSi0 3 , occurs in nature as the min- 
eral wollastonite, and, in combination with other silicates, 
in a large number of minerals, as garnet, mica, etc. It is 
formed when a solution of sodium silicate is added to a so- 
lution of calcium chloride, and when a mixture of calcium 
carbonate and quartz is heated to fusion. 

Glass. — Ordinary glass is a silicate of calcium and so- 
dium made by melting sand (silicon dioxide, silica, Si0 2 ) 
with lime and sodium carbonate or soda. Instead of cal- 
cium carbonate, lead oxide maybe used; and instead of 
sodium carbonate, potassium carbonate. The properties of 
the glass are dependent upon the materials used in its 
manufacture. 

Ordinary window-glass is a sodium-calcium glass. The 
purer the calcium carbonate and silica, the better the qual- 






GLASS. 315 

ity of the glass. This glass is comparatively easily acted 
upon by chemical substances, and is therefore not adapted 
to the preparation of vessels which are to be used to hold 
acids and other chemically active substances. It answers, 
however, very well for windows. The difference between 
ordinary window-glass and plate glass is essentially that 
the former is blown and then cut into pieces, while the 
latter, when in the molten condition, is run into flat 
moulds and there allowed to solidify. 

Bohemian glass is made with potassium carbonate. If 
pure carbonate is used, as well as pure calcium carbonate 
and silica, a very beautiful glass is the result. It is char- 
acterized by great hardness, by its difficult fusibility, and 
by its resistance to the action of chemical substances. It 
is particularly well adapted to the manufacture of vessels 
for use in chemical laboratories. 

Flint-glass is made by melting together lead oxide, po- 
tassium carbonate, and silicon dioxide. It is characterized 
by its power to refract light, its high specific gravity, its 
low melting-point, and the ease with which it is acted upon 
by reagents. Owing to its high refractive power, it is 
largely used in the manufacture of lenses for optical in- 
struments. 

Strass is a variety of lead-glass which is particularly 
rich in lead. Its refracting power is so great that it is 
used in the manufacture of artificial gems. 

Colors are given to glass by putting into the fused mass 
small quantities of various substances. Thus, a cobalt com- 
pound makes glass blue; copper and chromium make it 
green; one of the oxides of copper makes it red; ura- 
nium gives it a yellow color, etc. The most common 
variety of glass is that used in the manufacture of ordinary 
bottles. It is generally green to black, and sometimes 
brown. In its manufacture impure materials are used. 
chiefly ordinary sand, limestone, sodium sulphate, com- 
mon salt, clay, etc. 



316 INTRODUCTION TO CHEMISTRY. 

Glass which has been suddenly cooled is very brittle and 
breaks into small pieces when scratched or slightly broken 
in any way. This is shown by the so-called Prince Ru- 
pert's drops, which are made by dropping glass, in the 
molten condition, into water. When the end of such a drop 
is broken off, the entire mass is completely shattered into 
minute pieces. It is clear from this that, in the manufac- 
ture of glass objects, care must be taken not to cool them 
suddenly. In fact they are cooled very slowly, the process 
being known as annealing. For this purpose they are 
placed in furnaces the temperature of which is but little 
below that of fusion, and are kept there for some time, 
the heat being gradually lowered. If red-hot glass is in- 
troduced into heated oil or paraffin, and allowed to cool, 
it is found to be extremely hard and elastic. The glass 
of De la Bastie is made in this way. Vessels made of it 
can be thrown about upon hard objects without breaking, 
but sometimes a slight scratch will cause the glass to fly in 
pieces, as the Rupert's drops do. 

Mortar. — Mortar is made of slaked lime and sand. 
When this mixture is exposed to the air, carbonate of cal- 
cium is slowly formed, and the mass becomes extremely 
hard. The water contained in the mortar soon passes off, 
but nevertheless freshly-plastered rooms remain moist for 
a considerable time. This is due to the fact that a reac 
tion is constantly taking place between the carbon dioxide 
and calcium hydroxide by which calcium carbonate and 
water are formed, 

Ca(OH) 2 + 00 2 = CaC0 3 + H 2 0, 

and it is the water thus liberated that keeps the air moist. 
The complete conversion of the lime into carbonate re- 
quires a very long time, because the carbonate which is 
formed on the surface protects, to some extent, the lime in 
the interior. 



CEMENTS— CALCIUM SULPHIDE. 317 

It is generally regarded as unhealthy to live in rooms 
with freshly-plastered walls, because the air is constantly 
kept moist in consequence of the reaction above men- 
tioned. It is, however, difficult to see why the presence of 
a little extra moisture in the air should be unhealthy; and, 
if there is any danger from freshly-plastered walls, it 
seems probable that the cause must be sought for else- 
where. It is possible that the constant presence of moist- 
ure in the pores of the walls interferes with the important 
process of diffusion, and that therefore when the room is 
closed this natural method of ventilation cannot come into 
play. 

Cements. — When lime-stones which contain magnesium 
carbonate and aluminium silicate in considerable quanti- 
ties are heated for the preparation of lime, the product 
does not act with water as calcium oxide does, and this 
lime is not adapted to the preparation of ordinary mortar. 
On the other hand, it gradually becomes solid, in contact 
with water, for reasons which are not known. Such sub- 
stances are known as cements, or hydraulic cements. 
Other cements besides those made in the manner men- 
tioned are known. 

Calcium sulphide, CaS, is formed by heating calcium 
sulphate with charcoal. It is remarkable on account of 
the fact that it is phosphorescent. After having been ex- 
posed to sunlight, it continues to give light for some time 
afterward. This and the similar compound, barium sul- 
phide, are now used quite extensively in the preparation of 
luminous objects, such as match-boxes, clock-faces, plates 
for house-numbers, etc. 

Barium and Strontium. — The compounds of barium and 
strontium closely resemble those of calcium. Barium forms 
an oxide, BaO, corresponding to lime, and also another one 



318 INTRODUCTION TO CHEMISTRY. 

known as barium dioxide, Ba0 2 .* This is formed by pass- 
ing oxygen or air over barium oxide heated to a dull red 
heat. At a higher temperature it gives off the oxygen. 
These facts have been utilized for the purpose of extract- 
ing oxygen from the air. 

Barium oxide is converted into the hydroxide Ba(OH) 2 
when treated with water. This hydroxide is soluble in 
water, the solution being known as laryta-ivater. 

Flame Reactions. — Calcium compounds color the flame 
reddish yellow; strontium compounds, intense red; and 
barium compounds, yellowish green. 

Relations between the Atomic Weights of the Members 
of this Group. — Between the atomic weights of calcium, 
strontium, and barium there exists the same relation as 
that with which we are already familiar in other groups. 
The atomic weight of calcium is 40; of strontium, 87.5; 
and of barium, 137: 



40 + 137 _ 



88.5. 



*This compound lias already been referred to in describing the 
preparation of hydrogen dioxide, H 2 2 . When it is treated with 
sulphuric acid this reaction takes place : 

Ba0 2 + H 2 S0 4 = BaS0 4 + H 2 2 . 

When barium oxide, BaO, is treated with sulphuric acid this reac 
tion takes place : 

BaO + H 2 S0 4 = BaS0 4 + H 2 0. 

When barium dioxide is treated with hydrochloric acid, hydrogen 
dioxide is also formed thus : 

Ba0 2 + 2HC1 = BaCl 2 + H 2 2 . 

[Compare this with the action that takes place when hydrochloric 
ncid acts upon manganese dioxide.] 



CHAPTER XXII. 

THE MAGNESIUM GROUP : MAGNESIUM, ZINC, CADMIUM, 

Of the three members of this group, magnesium and 
zinc are by far the most common. 

Magnesium, Mg (At. Wt. 2-4). — Magnesium occurs very 
widely distributed in nature, and in considerable quantities. 
Among the important magnesium minerals are magnesite, 
which is the carbonate MgC0 3 ; dolomite, a double car- 
bonate of magnesium and calcium; soapstone, serpentine, 
and meerschaum, which is essentially a silicate of mag- 
nesium. Further, there are many silicates that contain 
magnesium, among them being asbestos and hornblende. 
The metal is also found in solution in many spring-waters 
in the form of the sulphate, or Epsom salt. 

Manufacture. — It is prepared by treating magnesium 
chloride with sodium at a high temperature. It is now 
manufactured in quantity by this method. 

Properties. — It is a silver-white metal with a high lustre. 
In the air it changes only slowly, but gradually becomes 
covered with a layer of the oxide. When heated above its 
melting-point in the air it burns with a bright flame, 
forming the white oxide. The light of the flame is very 
efficient in producing certain chemical changes, as the 
combination of hydrogen and chlorine. At ordinary tem- 
peratures magnesium does not decompose water; at 100° 

319 



320 



INTRODUCTION TO CHEMISTRY. 



it decomposes it slowly. [Note the marked difference in 
this respect between magnesium and the alkali metals.] 

Applications. —The principal use to which magnesium is 
put is for producing a bright light, as for photographing 
in spaces to which the sunlight does not have access, and 
for signalling. It is also used to some extent as an ingre- 
dient of materials employed in making fireworks. 

Compounds of Magnesium.— The chief compounds of 
magnesium are the oxide, MgO, called magnesia; the sul- 
phate, MgS0 4 + 7H 2 0, commonly called Epsom salt; 
the carbonate, MgC0 3 ; the silicates; and the chloride, 
MgCl,. 

Magnesium Oxide, MgO. — This compound is commonly 
called -magnesia. A fine white variety is made by heating 
precipitated magnesium carbonate; it is called magnesia 
usta. It is very difficultly soluble in water, forming with 
it magnesium hydroxide, Mg(OH) 2 , which is practically 
insoluble in water. [What difference is there between 
magnesium and calcium in this respect ?] 

Magnesium chloride, MgCl 2 , is of special interest for 
the reason that it is the compound from which the metal 
magnesium is made. It is prepared by dissolving the car- 
bonate in hydrochloric acid. On evaporating this solution 
to the proper concentration, crystals of magnesium chlo- 
ride containing water of crystallization, MgCl 2 + 6H 2 0, 
are deposited. When this compound is heated for the 
purpose of drying it, the larger part of it undergoes de- 
composition, thus : 

MgCl 2 + H 9 = MgO + 2HC1. 

The same thing takes place to some extent on heating 
calcium chloride with water, so that fused calcium chlo- 
ride is always slightly alkaline in consequence of the pres- 
ence of lime, or calcium oxide. 



MAGNESIUM SULPHATE-ZINC. 321 

Dry magnesium chloride is prepared by adding ammo- 
nium chloride to its solution and evaporating to dryness. 
A double chloride of the composition, MgCl 2 .NH 4 Cl, is 
formed which can be evaporated to complete dryness with- 
out decomposition. When perfectly dry this double salt 
breaks down at a high temperature into ammonium chloride 
and magnesium chloride. The ammonium chloride is 
volatilized, and the magnesium chloride remains behind. 

Magnesium Sulphate, MgS0 4 . — The mineral kieserite, 
which occurs at Stassfurt, has the composition MgS0 4 + 
H 2 0. The salt MgS0 4 + 7H 2 also occurs in nature. It 
is this variety which is generally obtained when a solution 
of magnesium sulphate is evaporated to crystallization. 
Its water solution has a bitter, salty taste. 

Uses. — Magnesium sulphate finds extensive application. 
It is used in medicine as a purgative, and is known as 
Epsom salt, as it is contained in the water of Epsom springs. 
It is used, further, in the manufacture of sodium sulphate 
and potassium sulphate, and as a fertilizer in place of 
gypsum. Its chief use is as a dressing for cotton goods. 



to, 



Zinc, Zn (At. Wt. G5). — Zinc, in almost all its com- 
pounds, exhibits a close resemblance to magnesium. It 
occurs in nature in combination principally as the carbonate, 
or smitlisonite, ZnCO s ; as the sulphide, or sphalerite, ZnS; 
and as the silicate, Zn 2 Si0 4 . Among other compounds of 
zinc found in nature are gahnite, Zn(A10 2 ) 2 , and frank- 
linite, which contains the compound Zn(Fe0 2 ) 2 with 
Fe(Fe0 2 ) a . 

Metallurgy. — The metallurgy of zinc is much simpler 
than that of magnesium, for the reason that the ores are 
easily converted into the oxide by roasting, and the oxide 
is easily reduced by heating it with charcoal. Owing to the 
volatility of the metal the vessels in which the reduction is 
effected must be so constructed as to facilitate the conden- 



322 INTRODUCTION TO CHEMISTRY. 

sation of the vapors. The vessels used are either earthen- 
ware muffles or tubes, open at one end and connected with 
iron receivers. At first the zinc vapor is condensed in the 
form of a fine dust, as in the case of sulphur. This forms 
the commercial product called zinc dust. It always contains 
zinc oxide. Afterwards the zinc condenses to the form of 
a liquid, and this is cast in plates. The zinc thus obtained 
is not pure, but contains lead and iron, and sometimes 
arsenic and cadmium. It is called spelter. By repeated 
distillation it can be obtained pure. When distilled under 
diminished pressure, it is deposited in beautiful lustrous 
crystals, the forms of which are extremely complicated. 

Properties. — Zinc has a bluish-white color and a high 
lustre. The crystals above referred to, which are perfectly 
pure zinc, have a brilliant lustre and do not change in the 
air. At different temperatures zinc has markedly different 
properties. At ordinary temperatures it is quite brittle; at 
100°-150° it can be rolled out in sheets, but above 200° it 
becomes brittle again. It melts at 433°, and boils at 1040°. 
When heated in the air it takes fire, and burns with a bluish 
flame, forming zinc oxide. This can be shown by means of 
the oxyhydrogen blowpipe. In dry air it does not change. 
Ordinary zinc dissolves in all the common acids, usually 
with evolution of hydrogen. In the case of nitric acid, 
however, the acid is to some extent reduced to ammonia. 
The purer the zinc the less readily is it acted upon by 
sulphuric acid, and the pure crystals above referred to are 
scarcely acted upon at all by this acid. Zinc also dissolves 
in the caustic alkalies, forming zincates. Pure zinc can be 
made to act upon sulphuric acid by adding a few drops of 
platinum chloride. 

Applications, — Zinc is extensively used as sheet-zinc, in 
making galvanic batteries, for galvanizing iron, etc. Zinc 
clust is a very efficient reducing agent, either in alkaline or 



ZINC IN ALLOYS-ZINC OXIDE. 323 

in acid solution. With caustic alkalies — as, for example, 
with potassium hydroxide — it gives hydrogen and a zincate : 

Zn + 2K0H - Zn(OK), + H 2 . 

With sulphuric acid also it gives hydrogen readily. Zinc 
is used in the preparation of important alloys. 

Alloys. — Iron covered with a layer of zinc is known as 
galvanized iron. As has been mentioned, zinc is a con- 
stituent of brass. It combines readily with mercury to form 
zinc amalgam, and this fact is taken advantage of for the 
purpose of preserving the zinc plates in galvanic batteries. 
Zinc plates covered with a layer of the amalgam are acted 
upon much more slowly than zinc itself. The amalgama- 
tion is effected by cleaning the zinc, dipping it in dilute 
sulphuric acid, and rubbing mercury over the surface with 
a brush or a piece of cloth. 

Zinc oxide, ZnO, is obtained as Flores zinci by burning 
zinc, and by heating the carbonate or nitrate of zinc. It 
turns yellow when heated, but on cooling becomes white 
again. 

Experiment 156. — Heat a small piece of zinc on charcoal in 
the oxidizing flame of the blowpipe. The white fumes of zinc 
oxide (philosopher's wool) will be seen, and the charcoal will be 
covered with a film which is yellow while hot, but becomes white 
on cooling. [What element gives a film which is white both when 
hot and when cold ?] 

Zinc oxide is used as a constituent of paint under the 
name of zinc white. 

Zinc sulphate, ZnS0 4 -f ?H 2 0, is commonly called white 
vitriol. [In w T hat experiments has zinc sulphate been ob- 
tained ?] It is obtained on the large scale by heating zinc 



324 INTRODUCTION TO CHEMISTRY. 

sulphide in contact with the air. Under these circum- 
stances, the sulphide is oxidized : 

ZnS + 40 = ZnS0 4 . 

This operation is known as roasting. By roasting zinc 
sulphide at a higher temperature it is converted into zinc 
oxide: 

ZnS + 30 = ZnO + S0 2 . 

Zinc sulphate is also formed in large quantities in gal- 
vanic batteries and in fche preparation of hydrogen. 

Zinc chloride, ZnCl 2 , is obtained by evaporating a water 
solution of the substance and distilling the residue. It is 
an oily liquid which has a very strong affinity for water. 
On evaporating a water solution a part of the chloride un- 
dergoes decomposition, just as magnesium chloride does, 
forming the oxide : 

ZnCl 2 + H 2 = ZnO + 2HC1. 

Some Insoluble Compounds of Zinc. — The hydroxide, sul- 
phide, carbonate, and phosphate of zinc are insoluble in 
water. 

Experiment 157. — Produce the insoluble compounds just men- 
tioned and express the reactions by means of equations. The 
phosphate of zinc precipitated by ordinary sodium phosphate is 
the normal phosphate, Zn 3 (P0 4 )2. 

[What happens on bringing together solutions of sodium car- 
bonate and zinc sulphate? ammonia and zinc chloride? barium 
chloride and zinc sulphate ? lime-water and zinc sulphate ? What 
color has zinc sulphide ? Is it thrown down when the solution 
contains dilute hydrochloric acid ? Try it.J 



CHAPTER XXIII. 

THE COPPER GROUP: COPPER, MERCURY, SILVER. 

Copper, Cu (At. Wt. 63.2). — Copper occurs in nature in 
the uncombined or native state in large quantities in the 
neighborhood of Lake Superior in the United States, and 
in China, Japan, Siberia, and Sweden. It also occurs in 
combination with oxygen as rubxj copper, which is the oxide 
Cu 2 0; with sulphur as chalcocite, Cu a S; and with sulphur 
and iron in copper pyrites, Cu 2 S.Fe 2 S 3 . 

Metallurgy. — It is obtained from the oxide by heating it 
with charcoal. [This reduction has been illustrated under 
the head of Carbon (see Experiment 89).] It is also obtained 
from the sulphides. The chemical changes involved are 
comparatively complicated. 

Properties. — Copper is a hard metal of a reddish color and 
metallic lustre. It does not change in dry air, but in moist 
air it gradually becomes covered with a green layer of a car- 
bonate of copper. Nitric acid dissolves it, copper nitrate, 
Cu(N0 3 ) 2 , being formed, and oxides of nitrogen evolved 
[explain the reaction]; hydrochloric acid does not act 
upon it; sulphuric acid acts when heated with the metal; 
the sulphate, CuS0 4 , is formed and sulphur dioxide given 
off [explain the reaction]. Copper does not decompose 
water, even when water-vapor is passed over the metal 
heated to red heat. [Compare with the conduct of the mem- 
bers of the potassium, calcium, and magnesium groups.] 

Dilute acids in general do not act upon it unless the air 

325 



326 INTRODUCTION TO CHEMISTRY. 

has access to it. This fact is of importance in connection 
with the use of copper vessels in culinary operations. Sub- 
stances containing vegetable acids can be boiled in bright 
copper vessels with impunity, for the water-vapor prevents 
the access of the air, but, on cooling, the air is admitted, 
and then action takes place, causing solution of some of 
the copper, which is objectionable. 

Precipitation of Copper. — Copper is precipitated from 
solutions of its salts by zinc, iron, and some other metals, 
and by an electric current. 

Experiment 158. — In a neutral solution of copper sulphate 
hang a strip of zinc. The zinc will become covered with a layer 
of copper, and zinc will pass into solution as zinc sulphate. The 
zinc simply displaces the copper in this case, as it displaces hydro- 
gen from sulphuric acid : 



Zn + CuS0 4 = ZnS0 4 -f- Cu ; 
Zn + H 2 S0 4 = ZnS0 4 + H 2 . 



Perform a similar experiment, using a bright strip of sheet- 
iron instead of the zinc. [What is the result?] To those who 
first performed this experiment the iron appeared to be changed 
to copper. [How would you go to work to determine whether the 
iron is changed to copper or not ?] 

Applications. — As is well known, copper is used very 
extensively for a variety of purposes, among which the fol- 
lowing may be mentioned : for electrical apparatus, coins, 
copper vessels, roofs, for covering the bottoms of ships, 
etc. It is also used in copper-plating; and in the prepara- 
tion of a number of valuable alloys, such as brass, bronze, 
gun-metal, bell-metal, etc. 

Alloys. — Brass is a mixture or compound of about one 
part of zinc and two parts of copper; these proportions 
may, however, be varied between quite wide limits. There 
is a variety of brass containing equal parts of zinc and cop- 
per, and another containing one part of zinc and five parts 






COMPOUNDS OF COPPER. 327 

of copper. Pinchbeck is made by combining two parts of 
copper and one of brass. 

Bronze consists of copper, zinc, and tin, The proportion 
of copper varies from 65 to 84 per cent; that of zinc from 
31.5 to 11 per cent; and that of tin from 2.5 to 4 per cent. 
When exposed to the air bronze becomes covered with a 
green coating of basic copper carbonate, which protects it 
from further action. This coating is now generally pro- 
duced artificially by a variety of methods, as by washing 
the surface with a solution of salts and acids. 

Gun-metal consists generally of copper and tin in the 
proportion of 11 parts of tin and 100 parts of copper. 

Bell-metal contains a larger proportion (from 20 to 25 
per cent) of tin than does gun-metal. 

Alloys with aluminium containing aluminium and cop- 
per in widely different proportions are made. That with 
3 per cent of copper has a w r hiter color than aluminium, 
the color being more like that of silver. On the other 
hand, an alloy of copper with 5 to 10 per cent of aluminium 
has a color resembling that of gold. This, which is known 
as aluminium bronze, is very hard and elastic, and is not 
easily acted upon by chemical reagents. It is now used to 
a considerable extent in the manufacture of ornamental 
and useful articles. 

German silver is an alloy consisting of copper, zinc, and 
nickel. 

Compounds of Copper. — Among the more common com- 
pounds of copper are the oxides Cu 2 and CuO; the sul- 
phate, CuS0 4 ; the carbonate; and the sulphide, CuS. 

Copper forms Two Series of Compounds. — Copper has the 
power to form two distinct series of compounds, of which 
the following are examples : 

CuCl, CuCl, ; 

CuBr, CuBr a ; 
Cu 2 0, CuO. 



328 INTRODUCTION TO CHEMISTRY. 

Those compounds which are of the first order, corre- 
sponding to the chloride CuCl, are called cuprous com- 
pounds. Thus, CuCl is cuprous chloride ; Cu 2 0, cuprous 
oxide, etc. On the other hand, compounds of the second 
order are called cupric compounds. Thus, CuCl 2 is cupric 
chloride; CuO, cupric oxide, etc. It has been suggested 
that perhaps the formula of the simpler cuprous com- 
pounds like CuCl, etc., should be doubled, and written 
Cu 2 Cl 2 , Cu 2 Br 2 , etc. This suggestion has its origin in the 
valence hypothesis. In cupric chloride, CuCl 2 , and cupric 
oxide, CuO, copper is evidently bivalent; whereas if the 
formulas of the cuprous compounds are the simple ones 
CuCl, CuBr, etc., then in them copper is univalent. If, 
however, cuprous chloride is Cu 2 Cl 2 , it may be that in it 
the copper is bivalent. It is only necessary to assume that 
in the molecule of cuprous chloride two atoms of copper 
are combined as represented thus : 



Cu— 

I 
Cu— 






If then each of the copper atoms should combine with a 
chlorine atom, the compound would have the formula 
Cu 2 Cl 2 . Unfortunately, we have no experimental means 
of showing whether the molecule of cuprous chloride is 
more probably Cu 2 Cl 2 or CuCl, so that the above reason- 
ing is purely speculative. It is better, therefore, for the 
present to keep to the simpler formula. Whatever the 
explanation may be, it is unquestionably a fact that there 
are two series of salts of copper, in one of which there is 
relatively half as much copper as in the other. Mercury, 
iron, and some other metals present similar phenomena. 



Cuprous oxide, Cu 2 0, is found in nature as ruby copper, 
and is formed when copper is heated in contact with the air. 
It is a bright-red substance insoluble in water. 






COMPOUNDS OF COPPER. 329 

Cupric oxide, CuO, is obtained by heating copper te> red- 
ness in contact with the air, or by heating the nitrate. It 
is also formed when caustic soda or potash is added to a 
boiling-hot solution of a copper salt. If the solution is 
cold, blue cupric hydroxide, Cu(OH) 2 , is precipitated, but 
this easily loses water, particularly if the solution is heated. 
The reactions which take place are : 

CuS0 4 + 2XaOH = Cu(OH) a + Na 2 S0 4 , and 
Ca(0H) 3 = CuO + H,0. 

Experiment 159. — Add some caustic soda or potash to a small 
quantity of a cold solution of copper sulphate in a test-tube. 
Heat and notice the change from blue to black. 

Copper Sulphate, CuS0 4 + 5H 2 0.— This salt is manufac- 
tured on a large scale and is commonly known by the name 
" blue vitriol/' [What salt is called " white vitriol " ?] It 
forms large blue crystals, and, when heated, loses water 
and becomes colorless. The colorless substance becomes 
blue again in contact with water. 

Copper sulphate is used extensively in the preparation of 
blue and green pigments, in copper-plating by electrolysis, 
in galvanic batteries, for preserving wood, etc. 

Copper sulphide, CuS, is a black substance which is 
formed by passing hydrogen sulphide through a solution of 
a copper salt, or by adding a soluble sulphide, as potassium 
sulphide or ammonium sulphide, to such a solution. 

Experiment 160. — Treat a dilute solution of copper sulphate 
with hydrogen sulphide, with ammonium sulphide, with potassium 
or sodium sulphide. 

Copper-plating. — The process of copper-plating consists 
in brief in depositing upon an object a layer of copper by 
putting it in a bath containing some copper salt and con- 
necting it with one pole of an electric battery. Decompo- 
sition of the copper salt takes place, and copper is deposited 



830 INTRODUCTION TO CHEMISTRY. 

upon the object. The process is extensively used in the 
preparation of electrotype plates. These are prepared 
either from wood-cuts or from type by making a mould of 
plaster of Paris, covering this with graphite, and immersing 
the mould thus prepared in the copper-plating bath. The 
plate thus made is an exact reproduction of the wood-cut 
or type of which the impression was taken. 

Mercury, Hg (At. Wt. 200). — Mercury occurs native as 
drops enclosed in rocks, though principally in combination 
with sulphur as cinnabar, HgS. It is obtained by roasting 
cinnabar, when vapors of mercury and sulphur dioxide are 
given off. The mercury is condensed in appropriate ves- 
sels. It is a silver-white metal of a high lustre. At ordi- 
nary temperatures it is liquid, though it becomes solid at 
— 39°. 5. Its specific gravity, water being the standard, is 
13.5959. It does not change in the air at ordinary temper- 
atures. It is insoluble in hydrochloric acid and cold sul- 
phuric acid. [Try each.] It dissolves in hot concentrated 
sulphuric acid, and is easily soluble in nitric acid. [Try 
each.] The vapor of mercury is very poisonous. 

Uses. — Mercury is extensively used in the manufacture 
of thermometers, barometers, etc. ; as tin amalgam for mir- 
rors; and in the processes by which gold and silver are 
obtained from their ores. 

Amalgams. — With other metals it forms alloys called 
amalgams. In ordinary galvanic batteries the zinc plates 
are treated with mercury, and thus covered with a layer of 
zinc amalgam which protects them from the action of the 
acids used. 

Compounds of Mercury. — Among the more common com- 
pounds of mercury are the oxide, HgO; the two chlorides, 
mercnrons chloride, HgCl, and mercuric chloride, HgCl q ; 



COMPOUNDS OF MERCURY. 331 

the two iodides, mercurous iodide, Hgl, and mercuric io- 
dide, Hgl 2 ; and the sulphide, HgS. 

Mercuric oxide, HgO, is the red substance which was used 
in one of our first experiments for the purpose of preparing 
oxygen. It was by heating this substance that oxygen was 
discovered, and the discovery of oxygen is perhaps the 
most important event in the history of chemistry. It is 
formed when mercury is heated for some time near its boil- 
ing-point in contact with the air, and is made by heating 
the nitrate. 

Mercurous chloride, HgCl, is commonly known by the 
name calomel. It is precipitated when a soluble chloride 
or hydrochloric acid is added to a solution of any mercu- 
rous salt. It is manufactured by subliming an intimate 
mixture of mercuric chloride and mercury : 

HgCl, + Hg = 2HgCl. 

It is a white substance, insoluble in water, which finds ex- 
tensive application in medicine. 

Mercuric chloride, HgCl 2 , commonly called corrosive sub- 
limate, is manufactured on the large scale by subliming an 
intimate mixture of mercuric sulphate and common salt: 

HgS0 4 + 2NaCl = Na 2 S0 4 + HgCl a . 

It is a white substance, soluble in water. It is extremely 
poisonous. It has a very marked influence upon the lower 
organisms that play so important a part in producing 
disease and the decay of organic substances. Wood impreg- 
nated with it is partly protected from decay. In surgery 
it is used for the purpose of preventing contamination of 
wounds by the hands and instruments of the surgeon. 

Mercuric sulphide, HgS, occurs in nature as cinnabar in 
the form of red crystals or crystalline masses. When pre- 



332 INTBODUCTION TO CHEMISTBY. 

pared artificially by rubbing mercury and flowers of sul- 
phur together, or by passing hydrogen sulphide through a 
solution containing a mercury salt, it is a black powder. 
When sublimed this powder yields red crystals. 

Precipitation of Mercury as Mercurous Chloride. — It will 
be noticed that of the two chlorides only mercurous chloride 
is insoluble in water. If any mercurous salt is present in a 
solution, mercurous chloride will be thrown down on add- 
ing a chloride or hydrochloric acid; whereas if the solution 
contains a mercuric salt the addition of a chloride or hy- 
drochloric acid will produce no precipitate. 

Silver, Ag (At. Wt. 108). — Silver occurs native; in com- 
bination with sulphur; and with sulphur and other metals. 
Small quantities of silver sulphide are found in almost all 
varieties of galenite or lead sulphide. It occurs more 
rarely as the chloride, bromide, and iodide. 

Metallurgy of Silver. — Much of the silver in use is ob- 
tained from galenite, PbS. This mineral is treated in 
such a way as to cause the separation of the lead (which 
see), and the silver is separated from sulphur at the same 
time. But it is dissolved in a large quantity of lead, and 
the problem which presents itself to the metallurgist is 
how to separate the small quantity of silver from the large 
quantity of lead. 

Pattinson's Method. — This is accomplished by melting 
the mixture and allowing it to cool until crystals appear. 
These are almost pure lead. They are dipped out by 
means of a sieve-like ladle, and the liquid left is again al- 
lowed to stand, when crystals are again formed, and these 
are removed in the same way as before. By this means, 
and by again melting the crystals removed, allowing the 
liquid to crystallize, and removing the crystals formed, 
there is finally obtained a product rich in silver, but which 



SILVER. 333 

still contains lead. This is heated in appropriate vessels 
in contact with the air, when the lead is oxidized, while 
the silver remains in the metallic state. This method of 
concentrating by crystallization of lead is known as Pat- 
tinson's method. 

Zinc Method. — Another method of separating lead and 
silver now 7 extensively used consists in treating the molten 
alloy with a small quantity of zinc. This takes up all the 
silver, and the alloy of zinc and silver thus formed is re- 
moved, and afterwards treated with superheated steam, by 
which the zinc is oxidized and the silver left unchanged. 

Amalgamation Process. — Some ores of silver are treated 
in another way, known as the amalgamation process. 
The ores are mixed with common salt and roasted, when 
the silver is obtained in the form of the chloride. This is 
then reduced to silver by means of iron and w T ater, the re- 
action taking place as represented in the following equa- 
tion: 

2AgCl + Fe = FeCl, + 2Ag. 

The mixture is next treated with mercury, which forms 
an amalgam with the silver, while the other metals pres- 
ent do not combine with the mercury. The amalgam can 
be separated from the rest of the mass without difficulty, 
and when heated to a sufficiently high temperature the 
mercury distils over, leaving the silver. 

Properties. — Silver is a white metal with a high lustre. 
It is not acted upon by air, oxygen, or water. Sulphur acts 
upon it readily, forming a black coating of silver sulphide. 
The metal is not dissolved by hydrochloric acid, but is 
easily dissolved by concentrated sulphuric acid and by 
dilute nitric acid. 



834 INTRODUCTION TO CHEMISTRY. 

Alloys of Silver. — The silver used for coins and most 
other purposes is an alloy with copper, the pure metal 
being too soft. The alloy usually contains from 7| to 10 
per cent of copper. Other metals covered with a layer of 
silver, deposited by the action of an electric battery, are said 
to be silver-plated. 

Compounds of Silver. — The principal compounds of silver 
are the chloride, AgCl; bromide, AgBr; iodide, Agi; and 
nitrate, AgN 3 . 

Silver nitrate, AgN0 3 , is known also by the name 
" lunar caustic." It is prepared by dissolving silver in di- 
lute nitric acid. 

Experiment 161.— Dissolve a 10- or a 25-cent piece in dilute 
nitric acid. [What action takes place ?] Dilute the solution 
to 200-300 cc. with water. [What is the color of the solu- 
tion ? What does this indicate? Does this color prove the 
presence of copper?] Add a solution of common salt until it 
ceases to produce a precipitate. [What is the chemical change?] 
Filter off the white silver chloride and carefully wash with hot 
water. Dry the precipitate on the filter, by placing the funnel 
with the filter and precipitate in an air-bath heated to about 110°. 
Remove the precipitate from the filter and put it into a porcelain 
crucible. Heat gently with a small flame until the chloride is 
melted. Cut out a piece of sheet-zinc large enough to cover the 
bottom of the crucible, and lay it on the silver chloride. Now 
acid a little water and a few drops of dilute sulphuric acid, and 
let the whole stand for tw 7 enty-four hours. The silver chloride is 
reduced to silver, and zinc chloride is formed : 

Zn + 2AgCl = ZnCl 2 + 2Ag. 

Take out the piece of zinc and wash the silver with a little 
dilute sulphuric acid, and then with water. Heat a small piece 
of the metal on charcoal w 7 ith the blowpipe flame until it melts 
and forms a bead. Dissolve the silver in dilute nitric acid and 
evaporate to dryness on the w 7 ater-bath, so that the excess of 
nitric acid is driven off. Dissolve the residue in w 7 ater and put 



SILVER IN PHOTOGRAPHY. 335 

the solution either in a bottle of dark glass or in one wrapped 
in dark paper. 

Experiment 162. — To a few cubic centimetres of water in a test- 
tube add 5 to 10 drops of the solution of silver nitrate just prepared. 
To this dilute solution add a little of a dilute solution of sodium 
chloride. The curdy white precipitate is silver chloride. Stand 
it aside where the light can shine upon it, and notice the change 
of color which gradually takes place. In the same way make the 
bromide by adding potassium bromide, and the iodide by adding 
potassium iodide, to silver nitrate. 

Application of Compounds of Silver in Photography. — 
It will be seen from the last experiments that the chloride, 
bromide, and iodide of silver are insoluble in water and 
are changed by light. The art of photography is based 
upon the changes which certain compounds, especially 
salts of silver, undergo when exposed to the light. Silver 
iodide is best adapted to most purposes. The salt is so 
changed by the light that when treated with certain com- 
pounds, such as ferrous sulphate, pyrogallic acid, etc., 
called u developers," a deposit of finely-divided silver is 
formed upon the plate in those places affected by the 
light. A plate of glass, or a sheet of properly prepared 
paper, is covered in the dark with a thin layer of a salt of 
silver. The plate is then exposed in the camera to the ac- 
tion of the light from some object to be photographed. 
The salt is changed where it is acted upon by the light, while 
where there is no light it is not acted upon. An image of 
the object towards which the plate was directed is thus 
left on the plate. But after the action of the developer is 
complete there is still upon the plate unchanged silver 
salt To remove this the plate is washed with a solution 
of sodium thiosulphate, Na 2 S 2 3 (hyposulphite), which dis- 
solves the salt. 

Precipitation of Metallic Silver. — Silver is precipitated 
from solutions of its salts by zinc, copper, mercury, and 
other metals. 



336 INTRODUCTION TO CHEMISTRY. 

Experiment 163. — To a solution of silver nitrate containing 
about 5 grams of the salt in 100 ec. water add a few drops of 
mercury, and let it stand. In a few days the silver will be de- 
posited in the form of delicate crystals. This formation is called 
the " silver- tree. 1 ' 

Insoluble Compounds of Silver.— The oxide, chloride, 
bromide, iodide, sulphide, "carbonate, and phosphate of 
silver are insoluble in water. 

Experiment 164.— Verify this statement. [What takes place 
when hydrochloric acid is added to a solution of a silver salt ? 
When silver nitrate is added to barium chloride ? When ammo- 
nium carbonate is added to silver nitrate ? When disodium phos- 
phate is added to silver nitrate ? In this case, normal silver 
phosphate, Ag 3 P0 4 , is formed, and some nitric acid is'set free.] 

Argentous and Argentic Compounds— Silver forms 
mostly those compounds which are analogous to the 
cuprous and mercurous salts, and not those which are analo- 
gous to the cupric and mercuric salts. There is, however, 
an oxide, Ag 2 0, and another, AgO, corresponding to mer- 
curous and. mercuric oxides. 

The Specific Heat of Elements as a Means of Determining 
their Atomic Weights. — The question naturally suggests 
itself, How are the atomic weights determined in the case 
of elements like silver, copper, etc., which cannot be con- 
verted into the form of vapor, and which do not yield com- 
pounds that can be converted into vapor ? It will be re- 
membered that most of the atomic weights with which we 
have thus far had to deal, as those of oxygen, chlorine, 
nitrogen, etc., are determined by a consideration of the 
specific gravity of the vapors of the compounds of these 
elements. The relative weights of equal volumes of these 
gases or vapors are determined, and then, assuming that 
these weights express the relative weights of the molecules 
of the compounds, the smallest weight of the element oc- 
curring in any compound is selected as the atomic weight. 
[Refer back and carefully read the chapter relating to the 



SPECIFIC IIEAT AND ATOMIC WEIGHT. 337 

Atomic Theory and Avogadro's Hypothesis.] But however 
valuable this method may be, it does not help us in the 
case of those elements which do not yield compounds capa- 
ble of conversion into vapor. In such cases the effect of 
heat upon the elements is of assistance. It has been found 
that when equal weights of different elements are exposed 
to exactly the same source of heat, they require different 
lengths of time to become heated to the same temperature. 
Given exactly the same heating power, a pound of water 
must be heated 32 times as long to raise its temperature 10, 
20, or 3j> degrees as a pound of mercury must be heated to 
raise its temperature the same number of degrees; or it takes 
32 times as much heat to raise a pound of water 10, 20, or 
30 degrees as it does to raise a pound of mercury the same 
number of degrees. The quantity of heat required to raise 
the temperature of a certain weight of a substance one 
degree as compared with the quantity of heat required to 
raise the temperature of the same weight of water one degree 
is called the specific heal; of the substance. Thus, from what 
was said above, the specific heat of mercury is ^, or, in 
decimals, 0.03332. In a similar way it can be shown that 
the specific heat of gold is 0.03244; of zinc, 0.0955; of sil- 
ver, 0.057; of copper, 0.0952. But these figures bear a re- 
markable relation to the atomic weights found by means 
of analysis. Thus, taking the above elements, we have: 

Specific Heat. At. Weight. 

Mercury 0.03332 200 

Gold 0.03244 196.2 

Zinc 0.0955 65 

Silver 0.057 108 

Copper 0.0952 63.2 

Calculation will show that the specific heat of these ele- 
ments is approximately inversely proportional to their com- 
bining weights. Thus, 

0.03332 : 0.057 :: 108 : 200 

Sp. Ht. of Hg. Sp. Ht. of Ag. At. Wt. of Ag. At. Wt. of Hg. 



338 INTRODUCTION TO CHEMISTRY. 

And the same is true in the other cases. Or the relation 
may be stated in another way, viz. : The product of the 
specific heat of any element multiplied by its combining 
weight is the same in all cases. The product is about 6.25. 
It is believed that the quantities of the elements to which 
this law refers are in reality the atomic weights, and we 
therefore accept the law known as the law of Dulong and 
Petit, which is this: 

The atomic weight of an element multiplied by its specific 
heat is a constant equal to about 6.25. 

There are some exceptions to the law, but these cannot 
be discussed at this time. Despite its imperfections it is 
now recognized as furnishing a valuable means of de- 
termining atomic weights. If A represents the atomic 
weight, and S the specific heat, then, according to the law 

6 25 
of Dulong and Petit, A X S = 6.25 nearly, and A = -W— 

To determine the atomic weight of an element by this 
method, then, it is only necessary to determine the specific 
heat of the element. Substituting for S the figure found, 
the value of A can be easily calculated. By careful analy- 
sis of compounds of the element the figure can be deter- 
mined more accurately. 



CHAPTER XXIV. 

THE ALUMINIUM GROUP: 

ALUMINIUM, GALLIUM, INDIUM, THALLIUM, SCANDIUM, 

YTTRIUM, LANTHANUM, AND YTTERBIUM. 

General. — The only element of this group which need be 
treated here is aluminium. This is an extremely impor- 
tant element that is found very widely distributed in nature. 

Aluminium, Al (At. Wt. 27). — Among the many impor- 
tant and widely distributed minerals which contain alu- 
minium are feldspar, granite, mica, and cryolite. 

Feldspar is a silicate of aluminium and potassium of the 
formula AlKSi 3 O s . Mica is a general name applied to a 
large number of minerals which are silicates of aluminium 
and some other metal, as potassium, lithium, magnesium, 
etc. The simplest form of mica is that represented by the 
formula KAlSi0 4 , according to which the mineral is a salt 
of orthosilicic acid, Si(OH) 4 . Cryolite is a double fluoride 
of aluminium aud sodium, or the sodium salt of fluoalu- 
minic acid, Xa 3 AlF 6 . Bauxite is a hydroxide of aluminium 
in combination with a hydroxide of iron. Aluminium 
occurs in the products of decomposition of minerals, as 
well as in the above forms. One of the most important of 
these is clay, which is found in all conditions of purity, 
from white kaolin to ordinary dark-colored clay. Kaolin 
is the aluminium salt of orthosilicic acid of the formula 
Al 4 (SiOJ, -f 4H 2 0. Aluminium silicate is found in all 
soils, but is not taken up by plants, and does not find 

339 



340 INTRODUCTION TO CHEMISTRY. 

entrance into the animal body. The name aluminium has 
its origin in the fact that the salt alum was known at an 
early date, and the metal was afterwards isolated from it. 

Preparation. — The preparation of aluminium on the 
large scale presents a problem of the highest importance 
to the human race. The element has properties which 
would appear to adapt it to most uses to which iron is 
put, and for many purposes it has advantages over iron. 
Further, we are supplied by nature with unlimited quanti- 
ties of the compounds of aluminium, which are distributed 
everywhere over the earth. While, however, iron, lead, tin, 
copper, and other metals can be isolated from their natural 
compounds without serious difficulty, aluminium, which is 
more abundant than any of them, and in some respects more 
valuable than any of them, is locked in its compounds so 
firmly that it is only by comparatively complicated and 
expensive methods that it can be isolated. 

The first method devised for the preparation of alumin- 
ium on the large scale consisted in heating aluminium 
chloride with sodium. The chloride was heated to boiling 
in a retort; the vapor passed through a vessel containing 
pieces of iron heated to redness, and then into a long tube 
containing sodium. Instead of aluminium chloride, the 
double chloride of aluminium and sodium, which is more 
easily prepared in the dry condition, is now used. The 
double chloride and cryolite are heated together with 
sodium in a properly-constructed furnace. It is, further, 
possible to prepare aluminium by electrolysis of the chlo- 
ride or of the double chloride above mentioned; and the 
oxide can be reduced by mixing it with charcoal and pass- 
ing the current from a powerful dynamo-machine through 
it. By the latter method an alloy of aluminium and copper 
is now prepared, but the preparation of aluminium alone 
by this method does not appear to be entirely successful. 
New methods for the preparation of the metal are con- 



ALUMINIUM. 341 

stantly being devised, and the price is constantly being 
lowered. The latest method of promise consists in the 
electrolysis of aluminium oxide, in the form of corundum, 
in a bath of molten cryolite contained in a carbon crucible. 
A large number of patents have been issued, covering 
methods for the preparation of aluminium; but these are 
frequently so imperfectly described, and the evidence of 
their value is so unsatisfactory, that it is difficult to pass 
judgment upon them. Until recently the commercial 
preparation of aluminium has appeared to be intimately 
connected with that of the commercial preparation of 
sodium; but, if the latest method is as good as is claimed, 
this is no longer the case. 

Properties. — The color of aluminium is like that of tin, 
and it has a high lustre. It is very strong, and yet malle- 
able. It is lighter than most metals in common use, its 
specific gravity being 2.5 to 2.7 according to the condition, 
while that of iron is 7.8, that of silver 10.57, that of tin 7.3, 
and that of lead 11.37. It does not change in dry or in 
moist air; and in the compact form it does not act upon 
water even at elevated temperatures. It melts at about 
700°, which is higher than the melting-point of zinc, and 
lower than that of silver. Hydrochloric acid dissolves it 
with ease, forming aluminium chloride. At ordinary tem- 
peratures nitric and sulphuric acids do not act upon it; 
at higher temperatures, however, action takes place, and 
the corresponding salts are formed. It dissolves in solu- 
tions of the caustic alkalies, forming the so-called alumi- 
nates. It reduces many oxides when heated with them to 
a high temperature; and it is used in the preparation of 
boron and silicon. 

Applications. — The metal is used to a considerable extent 
in the preparation of ornaments, and of useful articles in 
which lightness is of importance, as in telescopes and 



3-12 INTRODUCTION TO CHEMISTRY. 

opera-glasses. An alloy with a small percentage of silver 
is used for the beams of chemical balances. Aluminium 
bronze, which is an alloy with copper, is also used quite 
extensively. (See under Copper.) 

Compounds of Aluminium. — Among the more important 
compounds of aluminium are the oxide, A1 2 3 ; the hydrox- 
ide, Al(OH) 3 ; the alums; the silicates ; and the chloride, 
A1C1, 

Aluminium Oxide, A1 2 3 . — This compound occurs in 
nature in the form of ruby, sapphire, and corundum. It 
is very hard, and as emery is used for polishing. It is 
made artificially by heating the hydroxide, Al(OH) 3 : 

2A1(0H) 3 = A1 9 3 + 3H,0. 

Aluminium Hydroxide, Al(OH) 3 . — This compound is 
found in nature in crystallized form as hydrargillite. It is 
precipitated when ammonia is added to a solution of alu- 
minium sulphate: 

A1 2 (S0 4 ) 3 + 6NH 4 OH = 3(NH 4 ) 9 S0 4 + 2A1(0H) 3 . 

It forms a gelatinous mass which it is difficult to filter. 
[Precipitate some from a solution of ordinary alum.] The 
hydroxide is soluble in acids and in alkalies. In the 
former case salts are formed in which the hydroxide plays 
the part of a base; in the latter it acts like an acid. The 
salts formed with the alkalies are called alumi nates. In 
aluminium salts one atom of the metal replaces three 
atoms of hydrogen; thus, aluminium nitrate is A1(N0 3 ) 3 ; 
the sulphate, A1 2 (S0 4 ) 3 , etc. In the alu mi nates the three 
hydrogen atoms of the hydroxide are replaced by metal; 
thus, potassium alum mate is Al(OK) 3 , and sodium alumi- 
nate Al(ONa),. 



ALUMS-ALUMINIUM SILICATE. 343 

Experiment 165. — Precipitate some aluminium hydroxide from 
a dilute solution of alum, by means of caustic potash, and con- 
tinue to add the latter slowly, when the precipitate will dissolve. 
Do the same with caustic soda. 

Aluminium hydroxide, A1(0H) 3 , loses water when heated, 
and a compound of the formula A10 2 H is formed : 

Al(OH) 3 = A10 2 H + H 2 0. 

This compound is found in nature as the mineral dia- 
spore. It has acid properties and forms extremely stable 
salts, several of which are found in nature. Spinel is 
magnesium aluminate (A10 2 ) 2 Mg. The formation of the 
hydroxides Al(OH) 3 and A10 2 H, and of salts derived from 
each, indicates some analogy between aluminium and 
boron. On the other hand, the power to replace the hy- 
drogen of acids is not possessed by boron to any great ex- 
tent. [Refer back to Boron. Read again what is said 
about it, and compare it with aluminium.] 

Alums. — With the sulphates of the alkali metals alumin- 
ium sulphate forms complex compounds which crystallize 
beautifully. Potassium alum is the best known of these. 
It may be regarded as derived from 2 molecules of sul- 
phuric acid by the replacement of 3 atoms of hydrogen by 
1 atom of aluminium, and the fourth by 1 atom of potas- 
sium; thus, A1K(S0 4 ) 2 . The crystals always contain 12 
molecules of water, the complete formula being A1K(S0J 2 
+ 12H 2 0. Similarly, sodium alum is AlNa(S0 4 ) 2 + 
12H 2 0, and ammonium alum A1NH 4 (S0 4 ) 2 + 12H 2 0. 

Experiment 166. — Determine whether the alum in the labora- 
tory contains potassium or ammonium. Crystallize some. What 
forms do the crystals possess ? 

Aluminium Silicate. — The silicate of aluminium occurs 
in nature in enormous quantities. The most important of 
the minerals containing it are the feldspars, of which ordi- 



344 INTRODUCTION TO CHEMISTRY. 

nary feldspar, KAlSi 3 8 , is the most abundant. The feld- 
spars, again, enter into the composition of granite together 
with quartz and mica, and mica itself is a double silicate of 
aluminium, 
i 

Natural Decomposition of Feldspar. — Under the influence 
of moisture, the carbon dioxide of the air, and changes in 
temperature, the feldspars are undergoing slow decomposi- 
tion, the products being mainly potassium or sodium silicate 
and aluminium silicate. The salts of the alkali metals, 
principally the potassium salt, being soluble, are carried away, 
and find their way into the soil. The silicate of aluminium 
is not soluble, but it easily forms an emulsion with water, 
and is therefore carried down the sides of the hills and 
mountains upon which it is formed into the valleys, and 
much of it finds its way into streams. Sometimes this 
carrying away is prevented, and then beds of comparatively 
pure clay, known as kaolin, are formed. The clay found 
in the valleys is always more or less impure and colored. 

Kaolin. — This is the purest form of aluminium silicate 
found in nature. It always contains water. Its composi- 
tion varies, some specimens on analysis giving results which 
lead to the formula Al 4 (Si0 4 ) 3 + 4II 2 0, according to which 
the substance is the salt of normal silicic acid Si(OH) 4 . 
Other specimens have the composition HAlSi0 4 + H 2 0. 
When heated alone kaolin does not melt; but if feldspar 
is added to it, the whole melts, and forms a translucent 
mass known as porcelain. Other substances besides feld- 
spar may be used for this purpose. 

Clay. — Ordinary clay, as has been stated, is a name 
given to the impure varieties of aluminium silicate which 
have been carried down from the place of formation. 
Among the substances besides aluminium silicate found in 
days are calcium carbonate, magnesium carbonate, sand. 



ULTRAMARINE-PORCELAIN. 345 

and hydroxides of iron. The color is largely determined 
by the amount of the hydroxides of iron present. The 
better varieties are used in the manufacture of the so-called 
"stone-ware," gas-retorts, and fire-bricks. The colored 
varieties are used in making ordinary earthenware and 
bricks. Marl is clay mixed with considerable quantities of 
calcium carbonate. 

Ultramarine. — The substance occurring in nature and 
known as lapis lazuli consists of a silicate of sodium and 
aluminium together with a sulphur compound, probably a 
polysulphide of sodium. The coloring matter, known as 
ultramarine, obtained by ]30wdering it was formerly very 
expensive, but it is now made artificially by the ton, and 
the color of the artificially-prepared substance is even 
more beautiful than that of the natural. The artificial 
preparation is effected by melting together kaolin, an- 
hydrous sodium carbonate, and sulphur; or clay, calcined 
sodium sulphate, and charcoal. By varying the conditions 
of preparation, products of different colors are obtained. 
Besides the deep-blue ultramarine, there are now manufac- 
tured ultramarines of different shades of blue, and a green 
variety. 

Ultramarine is now manufactured in very large quan- 
tity — according to a recent report, to the extent of nearly 
9000 tons a year. It is the most extensively used blue 
coloring matter. 

Porcelain. — It was stated above that when kaolin is 
heated alone it does not melt, but that if feldspar is added 
to it, or if that found in nature contains feldspar, as is 
frequently the case, it either fuses together forming a com- 
pact mass, or melts and forms a translucent mass. Fur- 
ther, when kaolin or any other variety of clay is mixed 
with water, a plastic substance results, which can be 
kneaded and worked into any desired form. These facts 



346 INTRODUCTION TO CHEMISTRY. 

form the basis of the manufacture of earthenware, porce- 
lain, etc. The ease with which the mass melts depends 
upon the quantity of feldspar or other flux added to it. 
If but little is added it melts with difficulty; if much is 
added it melts easily. 

In the manufacture of the finest kinds of porcelain 
kaolin is used. This is generally mixed with a little feld- 
spar or chalk, gypsum, or some other flux, and sand is also 
added. All these substances must be very finely ground. 
The mixture is then worked into the desired forms, and 
carefully dried. After the objects are dried they are 
burned, first at a red heat at which the mass becomes solid, 
afterwards at a white heat for the purpose of forming a 
glaze upon the surface. The product after the first burn- 
ing is that which is familiar as porous earthenware; that 
formed in the second burning is the porcelain with glaze 
as it is commonly used. 

In order to form the glaze upon the porcelain the porous 
earthenware first formed is drawn through a vessel con- 
taining proper materials in finely-powdered condition and 
suspended in water. The materials used are generally the 
same as those used for the porcelain itself, but they are 
mixed in different proportions, with less kaolin and more 
sand and feldspar, so as to be more easily fusible. After 
this treatment the objects are again heated to a high tem- 
perature. 

Earthenware. — The ordinary varieties of earthenware 
are made from varieties of clay which are much less pure 
than kaolin. Ordinary colored clay is used. The ob- 
jects are formed, and then subjected in general to the 
same kind of treatment as porcelain. They are glazed in 
different ways. One method consists in bringing the glaz- 
ing material on the earthenware before it is burned; an- 
other method consists in putting the objects in the furnace 
without a glaze^ and towards the end of the firing process 



ALUMINIUM SALTS IN SOLUTION. 347 

sodium chloride is thrown into the furnace, and is thus 
brought in contact with the ware in the form of vapor. 
A chemical change takes place, resulting in the formation 
of a silicate of aluminium and sodium upon the surface. 
This melts, and forms a glaze. 

Bricks are the most common variety of unglazed earth- 
enware. Owing to the presence of other substances be- 
sides aluminium silicate, as, for example, calcium car- 
bonate, the material is comparatively easily fusible. The 
color of bricks is largely due to the presence of oxides of 
iron. 

Action of Soluble Carbonates and Soluble Sulphides on 
Solutions of Aluminium Salts.— With weak acids aluminium 
forms no salts. There is, for example, no carbonate. The 
sulphide is so unstable that it decomposes into the hy- 
droxide and hydrogen sulphide when exposed to moist air. 
When a soluble hydroxide is added to a solution of a salt 
of aluminium, the insoluble hydroxide is precipitated; but, 
as this has acid properties, it is dissolved in an excess of 
either caustic soda or caustic potash. Owing to the weak 
basic properties of the hydroxide, sodium carbonate and 
other soluble carbonates precipitate, not the carbonate, but 
the uncombined hydroxide. 

Experiment 167. — Add a dilute solution of sodium carbonate 
to a dilute solution of alum. The precipitate is the hydroxide : 

2A1K(S0 4 ) 2 4- 3Na 2 C0 3 4- 3H 2 

= K 2 S0 4 + 3Na 2 SO< + 3C0 2 + 2A1(0H) 3 . 

Filter off and, after washing carefully, show that the precipi- 
tate is not the carbonate. Try the same experiment with am- 
monium and potassium carbonates. 

When an aluminium salt, in solution is treated with am- 
monium sulphide, the hydroxide is precipitated. Even if 



348 INTRODUCTION TO CHEMISTRY. 

the sulphide were formed it would be decomposed into the 
hydroxide and hydrogen sulphide by water. 

Experiment 168. — Add ammonium sulphide to a solution of 
alum. The precipitate is aluminium hydroxide : 

2A1K(S0 4 ) 2 + 3(NH 4 ) 2 S + 6H a O 

= 3(NH 4 ) 2 S0 4 + K 2 S0 4 + 3H 2 S + 2A1(0H) 3 . 

Rare Elements of the Aluminium Group. — The other 
members of the aluminium group need not be taken up 
here. The existence and properties of two of them, gal- 
lium and scandium, were predicted by the aid of the 
periodic law, as has been pointed out (see p. 214). 



CHAPTER XXV. 

THE LEAD GROUP: LEAD, TIN, AND GERMANIUM. 

General. — The only two members of this group which 
need be studied here are lead and tin. There are some 
points of resemblance between them, but there are also 
marked differences. 

Lead, Pb (At. Wt, 20?). — Lead occurs in combination in 
several forms in nature; for example, as the sulphate, car- 
bonate, chromate, and sulphide. The sulphide, PbS, 
known as galenite, is the most important source of lead. 

Metallurgy. — The extraction of the metal from the sul- 
phide is accomplished in one of two ways: 

(1) By heating the sulphide with iron, when the latter 
combines with the sulphur, forming iron sulphide, while 
the lead is set free. 

(2) By roasting the sulphide until it is partly converted 
into lead oxide and lead sulphate, and then heating the 
mixture without access of air, when the following two re- 
actions take place: 

PbS + 2PbO =3Pb + S0 9 ; 

PbS + PbS0 4 = 2Pb + 2S0 2 . 

The lead is thus set free, and the sulphur is driven off 
as sulphur dioxide. 

349 



S50 INTRODUCTION TO CHEMISTRY, 

Properties. — Lead is a bluish-gray metal with a high 
lustre. It is very soft and not very strong. It melts at 
about 325°. All lead compounds are poisonous. Nitric 
acid dissolves it, but hydrochloric and dilute sulphuric 
acids do not. It is precipitated in metallic form from a 
solution of one of its salts by metallic zinc. The formation 
is sometimes called the " lead-tree" or " Arbor Saturni." 

Experiment 169. — Dissolve 5 grams lead nitrate* in a litre of 
water, and put the solution in a bottle. By means of a thread 
suspend a piece of granulated zinc in the centre of the solution, 
and let it stand. The lead will be deposited slowly in crystalline 
form. At the same time the zinc will pass into solution. The 
zinc simply replaces the lead : 

Zn + Pb(N0 3 ) a = Zn(N0 3 ) 2 + Pb. 

After the tree is formed, filter off some of the solution and see 
whether zinc is contained in it or not. There will probably be 
some lead left, so that in order to detect the zinc the lead will 
have to be removed first. This may be done by adding sulphuric 
acid and alcohol. The sulphate of lead is thus formed. As this 
is somewhat soluble in water and insoluble in alcohol, the latter 
is added. Filter off the lead sulphate, and to the filtrate add 
just enough ammonia to neutralize the sulphuric acid, and then 
ammonium sulphide. White zinc sulphide is precipitated. 

If all the lead is not precipitated by the sulphuric acid, the 
precipitate caused by ammonium sulphide will not be white, but 
more or less inclined towards black, according to the quantity of 
lead sulphide present. All the lead may be precipitated in the 
first instance by first adding some hydrochloric acid [What effect 
will this have on the solution of the lead salt ? Which chlorides 
are insoluble ?] and then passing hydrogen sulphide through the 
solution. Filter off and add ammonia and ammonium sulphide 
to the filtrate. 

Uses. — Lead is extensively used for a variety of purposes, 
as, for example, for sulphuric-acid chambers, for evaporat- 

* Instead of lead nitrate, the acetate, or sugar of lead, may be used. 



COMPOUNDS Off LEAD AND OXYGEN. 351 

ing-pans for alum and sulphuric acid, for shot, for water- 
pipes, and for making a number of valuable alloys. The 
use of lead water-pipes is a matter of much importance 
from a sanitary point of view. Ordinary drinking-water 
acts only very slightly upon lead, and not enough is dis- 
solved to be dangerous. Nevertheless, circumstances may 
at any time arise that will increase the solvent power of the 
water and serious results may follow. Air and water act 
together upon lead more readily than when the air is 
excluded. In moist air lead tarnishes. 

Experiment 170.— Cut a piece of sheet-lead an inch or two 
square and partly cover it with water in a shallow dish. Let it 
stand for several days, renewing the water from time to time. 
Bring all the washings together, filter, and examine the filtrate 
for lead. Test with (1) hydrochloric acid ; (2) sulphuric acid ; 
(3) hydrogen sulphide. Try the same experiments with a very 
dilute solution of lead nitrate or acetate. 

Compounds of Lead and Oxygen. — Lead forms three dis- 
tinct compounds with oxygen, viz. : lead suboxide, Pb 2 0; 
lead oxide, PbO; and lead peroxide, Pb0 2 . Red-lead, or 
minium,ha,s approximately the composition, Pb 3 4 , and is 
perhaps a mixture of the oxide and peroxide. 

Lead oxide, PbO, is known by the name of litharge. It 
is formed by the oxidation of molten lead in contact with 
the air. When litharge is heated in the air to 400° it takes 
up oxygen and is converted into minium, or red-lead, Pb 3 4 
(=: 2PbO + PbOj. When heated to a high temperature 
this gives up oxygen and is again converted into yellow 
lead oxide. Treated with nitric acid, a part is dissolved 
forming lead nitrate, while lead peroxide, a brown powder, 
remains behind. 

Experiment 171. — Treat a little minium with ordinary dilute 
nitric acid, and note the change in color. [Does lead pass into 
solution ? How do you know ?] 



352 INTRODUCTION TO CHEMISTRY. 

Lead peroxide, Pb0 2 , conducts itself somewhat like man- 
ganese dioxide. When treated with hydrochloric acid 
chlorine is evolved : 

Pb0 2 + 4HCi = PbCl 2 + 2H 2 + Cl 2 . 

Experiment 172.— Treat a little lead peroxide with concentrated 
hydrochloric acid in a test-tube. [In what form is the lead after 
the experiment ? Is the product soluble or insoluble in water?] 

Heated with sulphuric acid, oxygen is given off and lead 
sulphate is formed. 

Experiment 173. — Heat some lead peroxide with concentrated 
sulphuric acid. Show that oxygen is given off. [What is left be- 
hind ? Is it soluble or insoluble ?] 

Salts of Lead. — Among the more important salts of lead 
are the sulphate, PbS0 4 ; the nitrate, Pb(N0 3 ) 2 ; the car- 
bonate, PbC0 3 ; the acetate, Pb(C 2 H 3 2 ) 2 ; the chromate, 
PbCr0 4 ; and the sulphide, PbS. The acetate and nitrate 
are soluble in water; the others are not. 

Lead acetate, Pb(C 2 H 3 2 ) 2 , commonly called "sugar of 
lead," is the lead salt of acetic acid, C 2 H 4 2 , which is the 
acid contained in vinegar. It is formed by dissolving lith- 
arge in acetic acid. 

Insoluble Salts of Lead. — The sulphate, chromate, and 
chloride have already been referred to. They are formed 
by adding a soluble sulphate, chromate, and chloride to a 
solution of a lead salt. The chromate is the well-known 
chrome yellow. Lead chloride is soluble in hot water, but 
only slightly soluble in cold water. It crystallizes from its 
solution in hot water. 

Experiment 174. — To a dilute solution of lead nitrate or acetate 
add some hydrochloric acid. Heat and thus dissolve the precipi- 
tate, and stand it aside. On cooling, the lead chloride will crys- 
tallize out. It is not soluble in ammonia. [Does it differ from 
silver chloride in this respect ?] 



Salts of lead- tin. 353 

tead carbonate, PbC0 3 , occurs in nature as cerussite, and 
is precipitated by adding lead nitrate to a solution of am- 
monium carbonate; but when a solution of a lead salt is 
treated with a normal carbonate of sodium or potassium, 
a basic carbonate is precipitated. When, for example, an 
excess ot sodium carbonate is added to a solution of lead 
nitrate, the precipitate has the composition 3Pb0.2C0 2 + 
H 2 0. The salts usually obtained are more complicated 
than this. Basic lead carbonate is prepared and used 
extensively under the name of white-lead, as a pigment. 
An objecti/m to white-lead paint is that it turns dark under 
the influence of hydrogen sulphide. It also turns yellow 
in consequence of the action of some substance contained 
in the oil with which the lead carbonate is mixed. 

Lead sulphide, PbS, is the important mineral galenite 
to which reference has been made The compound is pre- 
cipitated by passing hydrogen sulphide through a solution 
of a lead salt, or by adding a soluble sulphide to such a 
solution. 

[How can you distinguish between a lead and a barium 
salt without using hydrogen sulphide ? Between lead and 
silver without using hydrogen sulphide or hydrochloric 
acid ? By hydrochloric acid alone ?] 

Tin, Sn (At. Wt. 118). — Tin occurs in nature mostly as 
stone, or cassiterite, which is the oxide Sn0 2 . 



m 



Metallurgy. — The ores are roasted for the purpose of 
getting rid of the sulphur and arsenic, and the oxide is 
then heated with coal in a furnace. After the reduction is 
complete the tin is drawn off and cast in bars. This tin is 
impure, and when again slowly melted, that which first 
melts is purer. By letting it run off as soon as it melts, the 
comparatively difficultly fusible alloy remains behind. 
The commercial variety known as Banca tin is the purest. 



354 INTRODUCTION TO CHEMISTRY. 

This is made at Banca in the East Indies. Block-tin, made 
in England, is also comparatively pure. 

Properties. — It is a white metal, which in general appear- 
ance resembles silver. It is soft and malleable, and can be 
hammered out into very thin sheets, forming thus the well- 
known tin-foil. At 200° it is brittle. It melts at 228°. 
It remains unchanged in the air at ordinary temperatures. 
It dissolves in hydrochloric acid, forming stannous chloride, 
SnCl 2 ; in sulphuric acid, forming stannous sulphate, 
SnS0 4 , sulphur dioxide being evolved at the same time. 
[Explain this.] Ordinary concentrated nitric acid oxidizes 
it, the product being a compound of tin, oxygen, and hy- 
drogen, known as metastannic acid, which is a white pow- 
der insoluble in nitric acid and in water. 

Uses. — It is used in making alloys, of which bronze, soft 
solder, and britannia metal are the most important. It is 
used most extensively for protecting other metals, as in the 
tinware vessels in such common use, which are made of 
sheet-iron covered with a layer of tin. 

Alloys. — Bronze has already been referred to under Cop- 
per. Soft solder is made of equal parts of tin and lead, or 
of two parts of tin and one of lead. Britannia metal is 
composed of nine parts of tin and one of antimony. Tin 
amalgam is made by bringing tin and mercury together, 
and is used in the silvering of mirrors. 

Stannous and Stannic Compounds. — Tin forms two classes 
of compounds, the stannous and stannic compounds. 
These do not bear to each other the same relation as that 
which exists between cuprous and cupric compounds. 
[What is this?] In stannous compounds the tin appears 
to be bivalent, as indicated by the formulas SnCl 2 , SnO, 
SnS, which respectively represent stannous chloride, oxide, 
and sulphide. In stannic compounds, on the other hand, 



STANNOUS AND STANNIC COMPOUNDS. 355 

the tin appears to be quadrivalent, as indicated by the for- 
mulas SnCl 4 , Sn0 2 , and SnS 2 , which respectively represent 
stannic chloride, oxide, and sulphide. 

In general, stannous compounds are readily converted 
into stannic compounds. 

Stannous chloride, SnCl 2 , is formed by dissolving tin in 
hydrochloric acid. If a solution of stannous chloride is 
added to a solution of mercuric chloride, or corrosive sub- 
limate, the latter is reduced to mercurous chloride, and 
this, being insoluble in water, appears as a precipitate. 
When stannous chloride and mercuric chloride are heated 
together in solution, metallic mercury is formed; 

2HgCl. 2 + SnCl, = 2HgCl + SnCl 4 ; 
HgCl a + SnCl, = Hg + SnCl, 

Experiment 1T5. —Dissolve a few grams of tin in ordinary 
dilute hydrochloric acid. Add a little of this solution to a solu- 
tion of mercuric chloride. A white precipitate of mercurous 
chloride will be formed. Heat the two solutions together, and 
notice the formation of metallic mercury, which appears as a 
gray powder 

Stannic oxide, Sn0 2 , occurs in nature as tin-stone. It is 
obtained by burning tin in the air. When melted together 
with caustic soda it dissolves as sodium stannate. This 
action suggests that which takes place when silicon dioxide 
is melted with an alkali and a silicate is formed, and when 
carbon dioxide and an alkali are brought together. The 
formulas of the products in these cases are similar, viz., 
Na a SnO s , Na 2 Si0 3 , and Na 2 CO s . 

Stannic Hydroxide, Sn(0H) 4 , and Stannic Acid, H 2 Sn0 3 . 
— Stannic hydroxide, Sn(OH) 4 , is perhaps formed when a 
solution of stannic chloride in water is boiled. The pre- 
cipitate obtained has, however, the composition H 2 Sn0 3 , and 
this is known as stannic acid. Stannic acid is precipitated 



356 INTRODUCTION TO CHEMISTRY 

also by treating a solution of a stannate with just enough 
of an acid to effect decomposition. The decomposition 
with hydrochloric acid takes place as represented in the 
equation 

Na 2 Sn0 3 + 2HC1 = 2NaCl + H 2 Sn0 3 . 

Metastannic Acid. — When tin is treated with concen- 
trated nitric acid it is converted into a white powder which 
is insoluble in water and in acids. Stannic acid is insoluble 
in water, but is easily soluble in hydrochloric, nitric, and 
sulphuric acids. The two acids cannot, therefore, be iden- 
tical though they appear to have the same composition. 
The product formed by oxidizing tin with nitric acid is 
called metastannic acid, 

Stannic chloride, SnCl 4 , is made by heating tin in 
chlorine. It is a heavy liquid, which distils at 120° without 
suffering decomposition. It fumes in the air very strongly. 
It has a marked affinity for water, and is used as a dehy- 
drating agent in a number of reactions. It has long been 
known by the name spiritus finnans Libavli* In solution 
it is obtained by dissolving tin in aqua regia. In this case 
the stannous chloride formed by the action of the hydro- 
chloric acid on the tin is oxidized by the nitric acid: 

SnCl 2 + 2HC1 + = SnCl 4 + H 2 0. 

Stannic sulphide, SnS 2 , is a yellow substance resembling 
arsenic sulphide. It is formed by passing hydrogen sul- 
phide through a dilute solution of stannic chloride. It is 
soluble in ammonium sulphide. 

Experiment 176. — Dissolve a little tin in aqna regia. Make 
the solution very dilute, and pass hydrogen sulphide through it. 
Filter off, wash, and treat with ammonium sulphide. [Does the 
precipitate dissolve ? Add an aeid to the solution. What takes 
place ?] 






HOW TO DISTINGUISH TIN. 357 

How to Distinguish Tin from Other Metals. — A peculiaritv 
of tin which distinguishes it from most other metals is its 
conduct towards nitric acid. As already stated, instead of 
dissolving in the acid, it is converted into a white, insoluble 
compound — metastannic acid. Antimony is also con- 
verted into a white oxide by nitric acid, but antimony does 
not dissolve in hydrochloric acid, while tin does. 

Experiment 177. — Treat a little tin with strong nitric acid, and 
notice the formation of the white metastannic acid. [Is it sol- 
uble in water?] Treat a little antimony in the same way. Now 
treat each element separately with hydrochloric acid. 

Experiment 178.— Examine a small piece of solder, and show 
that it contains lead and tin. Treat with aqua regia; dilute with 
water. [Will all the lead pass into solution under these circum- 
stances? Will any of it ?J Pass hydrogen sulphide through the 
much-diluted solution. Filter off the precipitate; wash with hot 
water ; treat with yellow ammonium sulphide ; filter ; add an 
acid to the filtrate. [Explain what takes place in each step.) 
The formation of a yellow precipitate, which is soluble in am- 
monium sulphide, is not conclusive evidence that tin is present, 
for arsenic sulphide has similar properties. In order to distin- 
guish between them advantage may be taken of the fact that 
arsenic sulphide is soluble in a solution of ammonium carbonate, 
while stannic sulphide is not. Treat some of the precipitate with 
a solution of ammonium carbonate ; filter ; add an acid, when, 
if any arsenic sulphide is in solution, it will be precipitated. 

Experiment 179.— Examine a small piece of bronze, and show 
that it consists of tin and copper. In this case, after getting the 
two metals in solution by means of aqua regia, dilute, and pass 
hydrogen sulphide through until the solution is saturated. 
Filter ; wash ; treat with yellow ammonium sulphide. Filter ; 
acidify ; prove that the yellow precipitate is not arsenic sulphide. 
Dissolve the black precipitate, which is mostly insoluble in ammo- 
nium sulphide, in nitric acid. [What change will the copper sul- 
phide undergo when treated with nitric acid?] Treat a little of 
the solution with caustic soda, and boil. [What changes take 
place ?] Filter and wash. Mix some of the black precipitate with 
sodium carbonate, and heat in the reducing flame of the blow- 
pipe. [What evidence of the presence of copper do you thus get ?] 



CHAPTER XXVI. 

THE IRON GROUP : IRON, COBALT, NICKEL. 

Iron, Fe (At. Wt. 5G). — At the present time it is un- 
doubtedly true that iron is the most important metal for 
man. It is not improbable, however, that in the future 
aluminium may take its place for many purposes, though 
there appears to be no immediate prospect of this inter- 
ference with the iron industry. 

Occurrence. — Iron occurs in the form of the oxides, Fe 3 4 
and Fe 2 3 ; as the carbonate, FeC0 3 ; in combination with 
sulphur as iron pyrites, pyrite, FeS 2 ; and as silicates and 
hydrated oxides, or hydroxides. The compounds princi- 
pally used in making iron are magnetite, Fe 3 4 ; haematite, 
Fe 2 3 ; brown iron ore, Fe 4 3 (OH) 6 ; and spathic iron, or 
siderite, FeC0 3 . 

Metallurgy. — After the ores are broken up, they are first 
roasted, in order to drive off water from the hydroxides; to 
decompose carbonates; to oxidize sulphides; and, as far as 
possible, to convert the oxides into ferric oxide, Fe 2 3 , 
which is the most easily reducible of the oxides of iron. 
After the ores are prepared in this w r ay they are reduced 
by heating them with carbon and fluxes in the blast- 
furnace, when the iron collects in the molten condition 
under the so-called slag at the bottom of the furnace. 
Blast-furnaces differ somewhat in construction, but the 
essential parts are represented in Fig. 54. 

356 



IRON. 



:m 



The inner cavity of the furnace is narrow at the top 
and bottom, as is shown in the figure. Through 
pipes, known as tuyeres, such as 
that represented at the lower part 
of the left-hand side of the figure, 
hot air is blown into the furnace 
to facilitate the combustion. In 
modern furnaces arrangements are 
made above for carrying off the 
gases and utilizing them as fuel. 
The inner walls are built of fire- 
bricks, and these are surrounded 
by ordinary bricks, or stone-work. 
The furnaces vary in height from 
25 to 80 or 90 feet, an average 
height being about 45 feet. The 
reduction of the ores is accom- 
plished by placing in the furnace 
alternating layers of coke or char- 
coal, and the ores mixed with 
proper fluxes. The nature of the 
flux depends upon the ore. If 
this contains silicon dioxide or clay, lime is added: 
while, if it contains considerable lime, minerals rich in 
silicic acid are used, such as feldspar, clay-slate, etc. 
The object of the flux is to form a slag in which the re- 
duced iron collects, and by which it is protected from 
oxidation. When the fire is once started in a blast-furnace 
the operation of reduction is continuous until the furnace 
is burned out. Alternate layers of ore and flux and carbon 
are added, and, as the reduced iron collects below, it is 
from time to time drawn off and allowed to solidify in 
moulds of sand. The operation requires close attention. 
The ores must be carefully studied, and the nature and 
amount of flux regulated according to the character of the 
ore, as above stated. Then, too, the temperature of the 




Fig. 54. 



360 INTRODUCTION TO CHEMISTRY. 

furnace is a matter of importance, and must be watched, 
and regulated by means of the blast. The reduction is 
largely accomplished by carbon monoxide. In the lower 
part of the furnace the fuel burns to carbon dioxide, but 
this comes in contact with hot carbon, and is then reduced 
to the monoxide. The hot monoxide in contact with the 
oxides of iron reduces these, and is itself converted into 
the dioxide. A large proportion of the carbon monoxide, 
however, escapes oxidation, and this is carried off from the 
top to the bottom of the furnace by properly-arranged 
pipes, and is then utilized as fuel. A furnace lasts from 
two to twenty years, and sometimes longer. 

Varieties of Iron. — The iron obtained as above described 
is known m pig-iron or cast-iron. It is very impure, con- 
taining carbon, phosphorus, sulphur, silicon, etc. If, when 
drawn from the furnace, the iron is cooled rapidly, nearly 
all the carbon contained in it remains in chemical combi- 
nation, and the iron has a silver-white color. This product 
is known as white cast-iron. If the iron cools slowly, most 
of the carbon separates as graphite, and this being distrib- 
uted through the mass gives it a gray color. This product 
is known as gray cast-iron. If the ore contains consider- 
able manganese, this is reduced with the iron, and iron 
made from such ores and containing manganese has the 
power to take up more carbon than ordinary iron. This 
product, containing from 3.5 to 6 per cent combined car- 
bon, is known as spicgel-iron. 

All varieties of cast-iron are brittle, and easily fusible. 
The gray iron fuses at a lower temperature than the white, 
and is not as brittle; it is therefore well adapted to making 
castings. When cast-iron is treated with hydrochloric acid 
the carbon which is present in combined form is given off 
in combination with hydrogen as hydrocarbons, some of 
which have a disagreeable odor. This is, of course, the 
cause of the bad odor noticed on dissolving ordinary cast- 



IRON- STEEL. 361 

iron in acids. The uncombined or graphitic carbon, on 
the other hand, remains undissolved. Owing to its brittle- 
ness, cast-iron cannot be welded. When the carbon, silicon, 
and phosphorus are removed the iron becomes tough and 
malleable, and its melting-point is much raised. The prod- 
uct thus obtained is known as wrought-iron. 

Cast-iron is converted into wrought-iron in one of two 
ways : 

(1) By melting it in contact with the air, and blowing 
air into the molten mass. The carbon, phosphorus, and 
silicon are thus oxidized and got rid of. This process is 
known as puddling. 

(2) By mixing cast-iron with some of the purer ores and 
heating to a high temperature, when the carbon, phosphorus, 
etc., are oxidized by the oxygen of the ores. This process 
is called cementation. 

Wrought-iron contains less than 0.6 per cent of carbon, 
and, as the percentage of carbon decreases, the malleability 
increases and the melting-point rises. The melting-point 
of good wrought-iron is from 1900° to 2100°. Small 
quantities of sulphur, phosphorus, silicon, and manganese 
exert a very marked influence upon its properties. The 
process of welding consists in heating two pieces of iron to 
a high temperature, putting some sand upon one of them, 
laying them together, and hammering, when, as is well 
known, they adhere firmly together. The object of the 
sand is to keep the surfaces bright, which it does by unit- 
ing with the oxide and forming an easily fusible silicate. 

Steel. — There are two general methods of making steel : 

(1) W r rought-iron is heated with charcoal or with iron 
containing carbon in such proportions that the entire mass 
contains the necessary percentage of carbon to form steel. 
This is known as the cementation process. 

(2) Cast-iron is melted in a large vessel called a converter, 
Fig. 55, and then partly oxidized by means of blasts of air 



362 



INTRODUCTION TO CHEMISTRY. 




forced into the mass. Cast-iron is now added, and steel 
containing any desired percentage of carbon thus made. 
This is known as the Bessemer process. 

The Bessemer process is now employed on an enormous 
scale, there being a large demand for 
the product. It is used in making 
cannon, rails, axles, etc. Iron contain- 
ing more than a very small percentage 
of phosphorus is not adapted to the 
manufacture of Bessemer steel in the 
ordinary way; but it has been found 
that, if the converters are lined with 
lime-stone, such iron may be used. 
Under these circumstances the phos- 
phorus is oxidized, and with the lime- 
stone forms calcium phosphate, which 
is of value as a fertilizer (see Calcium 
Phosphate). This process is known as 
the Thomas-Gilchrist or the basic-lining process. 

Cast-iron can, further, be converted into steel by heating 
it together with pure iron ores. This method is employed 
in making Uchatius steel and, more extensively, Martin 
steel. 

Steel contains less than 1.5 per cent combined carbon, 
has a grayish-white color, and is capable of a high polish. 
When a steel rich in carbon is heated and cooled suddenly, 
it is rendered extremely hard and elastic; and when 
hardened steel is carefully heated, and allowed to cool 
slowly, it becomes very tough. This treatment is called 
tempering. 

Uses. — Any statement in regard to the applications of 
the three varieties of iron, viz., cast-iron, steel, and wrought- 
iron, would be superfluous. 



Fig. 55. 



Properties of Iron.— Pure iron is almost unknown. Of 



PROPERTIES OF IRON. 363 

the commercial varieties, it follows from what has been 
said that wrought-iron is the purest. That which is used 
for piano-strings is the purest commercial iron; it contains 
only about 0.3 per cent of impurities. Pure iron can be 
made in the laboratory by igniting the oxide or oxalate in a 
current of hydrogen, and by reducing ferrous chloride in 
hydrogen. In larger quantity it can be prepared by melt- 
ing the purest wrought-iron in a lime crucible by means of 
the oxhydrogen flame. The impurities are taken up by the 
crucible, and a regulus of the pure metal is left behind. 
That made by reduction of the oxide or oxalate is, of 
course, in finely-divided condition. If in its preparation 
the temperature is kept as low as possible, the product 
takes fire when brought in contact with the air; while if 
the temperature is high, the product has not this power. 
Pare iron is white and is one of the hardest metals. Its 
melting-point is higher than that of wrought-iron. Pure 
iron is attracted by the magnet. In contact with a magnet, 
or when placed in a coil through which an electric current 
is passing, it becomes magnetic; but the purer it is the 
sooner it loses the magnetic power when removed from 
the magnet or the coil. Steel, however, retains its mag- 
netism. When heated to a sufficiently high temperature 
iron burns, and forms the oxide, Fe 3 4 . This takesplace 
much more easily in oxygen than in the air. In dry air 
iron does not undergo change, but in moist air it rusts, 
or it becomes covered with a layer of oxide and hydrox- 
ide, which is formed by the action of the air, carbon 
dioxide, and water. The presence of salts in solution 
facilitates the rusting. Various methods are adopted to 
protect iron from this change, most of which are, how r ever, 
purely mechanical. A method that promises valuable re- 
sults is that invented by Barff. This consists in introduc- 
ing the iron into water-vapor at a temperature of C50°, 
when it becomes covered with a firmly-adhering layer of 
oxide. 



364 INTRODUCTION TO CHEMISTRY. 

Iron dissolves in acids with evolution of hydrogen, and 
generally with formation of ferrous salts: 

Fe + 2HC1 = FeCl 2 + H 2 ; 
Fe + II 2 S0 4 = FeS0 4 + H 2 . 

When cold nitric acid is used, ferrous nitrate and ammo- 
nium nitrate are the products; if the acid is warmed, ferric 
nitrate and oxides of nitrogen are formed. When an iron 
wire which has been carefully polished is introduced for an 
instant into red fuming nitric acid it can afterward be put 
into ordinary nitric acid without undergoing change. It is 
said to be in the passive state; and the commonly-accepted 
explanation of the phenomenon is that the wire is covered 
with a thin layer of oxide. As, however, the passive con- 
dition is lost by contact with an ordinary wire, the expla- 
nation does not appear to be adequate. 

Iron forms Two Series of Compounds. — Iron, like mer- 
cury, copper, and tin, forms two series of compounds that 
differ markedly from each other. These are the ferrous 
and ferric compounds. Thus with chlorine it forms two 
chlorides, one of which, ferrous chloride, has the composi- 
tion expressed by the formula FeCl 2 ; the other, ferric 
chloride, by FeCl 3 . It appears from a study of the specific 
gravities of the vapors of these chlorides that the above 
formulas should be doubled, so that ferrous chloride is now 
commonly represented by Fe 2 Cl 4 , and ferric chloride by 
Fe,Cl.. 

Similarly there are two oxides, FeO and Fe 2 3 ; two sul- 
phates, ferrous sulphate, FeS0 4 , and ferric sulphate, 
Fe a (S0 4 )„etc. 

Ferrous Compounds are converted into Ferric Compounds 
by Oxidation. — Ferrous compounds show a tendency to pass 
into ferric compounds by simple contact with the air; and 
are readily converted by oxidizing agents, such as nitric 



FERROUS AND FERRIC COMPOUNDS 365 

acid, potassium chlorate, etc. When, for example, ferrous 
hydroxide, Fe(OH) 2 ,* is exposed to the air suspended in 
water, it changes to ferric hydroxide, Fe(OH) 3 . The 
change is represented by the equation 

2Fe(OII) 2 + H 2 + = 2Fe(OH) 3 . 

So, also, when ferrous chloride is left standing in hydro- 
chloric-acid solution it changes to ferric chloride, and the 
change is rapidly effected by boiling with a little nitric acid : 

2FeCl 2 + 2HC1 + = 2FeCl 3 + H 2 0. 

Ferrous chloride, FeCl 2 , is formed by dissolving iron in 
hydrochloric acid. 

Experiment 180.— Dissolve a little iron wire in dilute hydro- 
chloric acid. Hydrogen is evolved, accompanied by small quan- 
tities of other gases, the formation of which is due to the presence 
of impurities in the iron, and carbon is left undissolved as a black 
residue. To a little of the solution in a test-tube add at once 
caustic soda. This precipitates ferrous hydroxide, Fe(OH) 2 , 
which changes color rapidly, becoming reddish-brown finally. 
Pure ferrous hydroxide is white. As it passes to the ferric con- 
dition it becomes dirty green, and darker and darker until it is 
reddish brown. Heat another portion of the solution of ferrous 
chloride to boiling, add two or three drops of concentrated nitric 
acid, and boil again. Repeat this operation two or three times. 
The ferrous chloride is thus oxidized to ferric chloride. It will 
be noticed that the color of the solution after the oxidation is 
reddish yellow, whereas before the oxidation it was nearly color- 
less or greenish. Add caustic soda to the solution of ferric chlo- 
ride. A reddish-brown precipitate of ferric hydroxide will be 
formed. Just as in this case we have passed from ferrous chlo- 

* If ferrous chloride lias the formula Fe 2 Cl 4 , it seems probable 
that the formula of ferrous hydroxide is Fe 2 (OH) 4 . We have no evi- 
dence in regard to this, and hence the simpler formula may be used 
here, particularly as we are for the present interested mainly in the 
composition of the compound. 



366 INTRODUCTION TO CHEMISTRY. 

ride to ferric chloride by oxidation, so we can pass back again to 
the ferrous compound. Thus, by adding a little zinc to a solu- 
tion of ferric chloride in which hydrochloric acid is present, the 
hydrogen evolved extracts chlorine from the ferric chloride and 
converts it into ferrous chloride : 

FeCl 3 + H = FeCl 2 + HC1. 

Ferrous Sulphate, FeS0 4 + 711,0.— This salt, which is 
commonly known as " green vitriol " or " copperas/' is 
formed by the action of sulphuric acid on iron. [What is 
"white vitriol," "blue vitriol"?] It undergoes change in 
the air, being converted into a compound containing ferric 
sulphate, Fe 2 (S0 4 ) 3 , and ferric hydroxide: 

GFeS0 4 + 30 + 3H 9 = 2Fe 2 (S0 4 ) 8 + 2Fe(0H) :H . 

Iron alum, FeK(S0 4 ), + 12H B 0,.is formed by bringing 
ferric sulphate and potassium sulphate together. It re- 
sembles ordinary alum, A1K(S0 4 ) 2 + 12H 2 0, but differs 
from it in containing iron instead of aluminium. 

Ferrous oxide, FeO, cannot be prepared in pure condi- 
tion on account of its great affinity for oxygen. 

Ferric oxide, Fe 2 C 3 , occurs in nature in lustrous crys- 
tals as hmmatite, and in other valuable ores of iron. The 
hydroxide corresponding to this — viz., ferric hydroxide, 
Fe(OH) 3 — is analogous in composition and properties to 
aluminium hydroxide. It is a weak base, but, unlike alu- 
minium hydroxide, it does not form compounds with bases. 
Hence it does not dissolve in caustic soda and caustic potash. 
[Try it. Suppose a solution contains an aluminium salt 
and a ferric salt, and caustic soda is added, what will first 
take place? If more is added and the solution filtered, 
where will the aluminium be found, and where the iron ?] 

Certain natural compounds of iron appear to be deriva- 
tives of the hydroxide FeO. OH, which corresponds to the 



FERROUS AND FERRfC COMPOUNDS-NICKEL. 367 

aluminium compound A10.0II, and to metaboric acid, 
BO. OIL Thus magnetite is thought to be the ferrous salt 

of this acid as represented by the formula ,, .'^ > Fe; 

and franklinite is probably the corresponding zinc salt 

FeO.O ~ 

FeO.O >Zu - 

Ferroso-ferric oxide, Fe 3 4 , or magnetic oxide of iron, is 
found in nature in the form of loadstone. It is formed 
when iron is burned in oxygen (see Experiment 24). 

Ferric Acid, lI. 2 Fe0 4 . — It is interesting to note that iron 
combines with a larger proportion of oxygen than is con- 
tained in any of the compounds thus far mentioned, and 
then forms an acid. Potassium ferrate has the composition 
represented by the formula K 2 Fe0 4 . 

Sulphides of Iron. — The sulphides of iron have been re- 
peatedly mentioned. Ferrous sulphide, FeS, is made by 
heating sulphur and iron together in proper proportions. 
It is used in making hydrogen sulphide [Explain how]. 

Iron pyrites, or pyrite, FeS 2 , is a yellow crystallized sub- 
stance found very abundant in nature. When heated in 
a closed tube, sulphur is given off. When heated in an 
open vessel, as upon a shallow iron pan, the sulphur is oxi- 
dized to sulphur dioxide, and the iron is left in the form 
of the oxide. [Verify these statements by experiment.] 

Nickel, Ni (At. Wt. 58. G), is found in meteoric iron and 
in combination with arsenic. It forms two series of salts 
corresponding to the two hydroxides nickelous h} r droxide, 
Ni(OH) 2 , and nickelic hydroxide, Ni(OH) 3 . 

Most nickel salts are colored green. 

Alloys of nickel are extensively used. Argentan or Ger« 



368 INTRODUCTION TO CHEMISTRY. 

man silver consists of copper, zinc, and nickel. Various 
nickel alloys are used for making coins. The 5- and 3-cent 
pieces of the United States are made of an alloy consisting 
of 25 per cent nickel and 75 per cent copper. 

Nickel is, further, extensively used in electro-plating. 
Iron is covered with a thin layer of the metal to protect 
it from rusting. 

Cobalt, Co (At. Wt. 58.7), is found in combination with 
arsenic and sulphur, and also in small quantities accom- 
panying nickel in meteoric iron. 

The salts of cobalt are red in combination with water, 
and blue when dried. If marks are made on paper with a 
dilute solution of one of the salts the color is not percepti- 
ble. If, however, the paper is held before a fire, the salt- 
loses water and turns blue, and, as the blue is more in- 
tense than the red, it is visible. When the salt is again 
moistened it becomes invisible. This is the basis of the 
preparation of the so-called sympathetic inks. 



CHAPTER XXVII. 

MANGANESE.— CHROMIUM.— URANIUM. 

Manganese, Mn (At. Wt. 55). — Manganese is found in 
nature in the form of oxides, of which manganese dioxide, 
or the black oxide of manganese, occurs most frequently. 

Compounds of Manganese with Oxygen. — With oxygen it 
forms the following compounds: manganous oxide, MnO; 
manganic oxide, Mn 2 3 ; manganoso -manganic oxide, 
Mn 3 4 ; manganese dioxide, Mn0 2 ; and permanganic an- 
hydride, Mn.,0 7 . 

Comparison of Manganese with Aluminium and with 
Iron. — Manganese presents points of resemblance with 
aluminium and iron. Like iron it forms two series of 
salts, the manganous and manganic series, which differ 
from each other very much as ferrous and ferric salts do. 
Like iron, also, it forms an oxide, Mn 3 4 , which is analo- 
gous to the magnetic oxide of iron. L'nlike iron, it forms 
the dioxide MnO.,. Like iron, it forms salts, which are de- 
rived from an acid of the formula H,Mn0 4 ; as, for exam- 
ple, potassium manganate,T£ 2 M.nO A . Unlike iron, it forms 
salts derived from an acid HMn0 4 ; as, for example, potas- 
sium permanganate, KMn0 4 . 



Formation of Manganous Salts. — All the higher oxides of 
manganese lose a part of their oxygen very easily, and 
are usually converted into manganous salts, like MnS0 4 , 
MnCl 2 , etc., in which the metal is apparently bivalent. 

369 



S70 INTRODUCTION TO CHEMISTRY. 

The use of manganese dioxide in preparing oxygen and 
chlorine has been described. [Give an account of the 
changes which manganese dioxide undergoes when treated 
with sulphuric acid; hydrochloric acid; when heated.] 

Manganese Dioxide, Mn0 2 . — This important compound 
occurs in nature in considerable quantity, and is known as 
pyrolusite or black oxide of manganese. The chief appli- 
cation of the dioxide is in the preparation of chlorine. It 
is also used for making oxygen, and for the purpose of <fle- 
colorizing glass. In the last process a small quantity is 
added to the molten glass. This alone would give the 
glass an amethyst color. Without it the glass would be 
green. One color counteracts the other, and the glass 
appears colorless. 

Weldon's Process for the Regeneration of Manganese Di- 
oxide in the Preparation of Chlorine. — In the preparation 
of chlorine by means of manganese dioxide and hydro- 
chloric acid the manganese is converted into manganous 
chloride which is practically worthless. Wei don has, how- 
ever, devised a method for the utilization of the waste- 
liquors of the chlorine factories. When manganous chlo- 
ride in solution is treated with lime the corresponding 
hydroxide is precipitated : 

MnCl, + Ca(OH) 2 = Mn(OH) 2 + CaCl 2 ; 

and when this hydroxide mixed with lime is treated with 
steam and air oxidation takes place, and a compound 
CaMn0 3 or CaMn a 6 is formed: 

Mn(0H) o + Ca(0H) 2 + = CaMnO, + 2H 2 0; 
2Mn(OH) 2 + Ca(OH) 2 + 20 = CaMn 2 5 + 3H 2 0. 

These compounds give chlorine when treated with hydro- 
chloric acid. They may indeed be regarded as consisting of 
lime and manganese dioxide, CaO.Mn0 2 and Ca0.2Mn0 2 . 



POTASSIUM PERMANGANATE. 371 

Potassium Permanganate, KMn0 4 . — This salt is obtained 
from potassium manganate, K 2 Mn0 4 , by boiling or by 
passing carbon dioxide into it. The manganate is made 
by treating manganese dioxide with potassium hydroxide 
and potassium chlorate; in other words, by oxidizing man- 
ganese dioxide in the presence of the base, potassium hy- 
droxide. The reaction is represented by the equation 

3MnO f + GKOH + KC10 3 = 3K 2 Mn0 4 + KC1 + 3H 4 0. 

The permanganate is a dark-colored, crystallized com- 
pound which dissolves in water, forming a deep purple- 
colored solution. 

Experiment 181. — In a small porcelain crucible heat together 
5 grams powdered manganese dioxide, MiiOa , 5 grams solid potas- 
sium hydroxide, and 2£ grams potassium chlorate, KC10 3 . When 
the mass has turned green, let the crucible cool, dissolve the con- 
tents in water and boil the solution. The green substance is po- 
tassium manganate. The color will change from green to purple. 

Reduction cf Potassium Permanganate. — Potassium per- 
manganate gives up its oxygen very readily and changes 
to a hydroxide of manganese. If an acid is present the 
hydroxide dissolves, forming a colorless solution. When, 
therefore, a solution of potassium permanganate is added 
to an acid solution containing an oxidizable substance it 
becomes colorless. 

Experiment 182. — To a dilute solution of ferrous sulphate 
containing free sulphuric acid add drop by drop a dilute solution 
of potassium permanganate. The color will be destroyed as long 
as there is any ferrous sulphate present. 

Add some permanganate solution to a solution of sulphur di- 
oxide in water. [What would you expect to take place in this 
case ?] 

Add some comparatively dilute hydrochloric acid to a few crys- 
tals of potassium permanganate in a test tube. [What do you' 
notice ? How do you explain the change ?] 



372 INTRODUCTION TO CHEMISTRY. 

Comparison of Potassium Permanganate with Potassium 
Perchlorate. — Potassium permanganate, KMn0 4 , is analo- 
gous to potassium perchlorate, KC10 4 , not only in com- 
position, but in its general properties. 

Chromium, Cr (At. Wt. 52.3). — This element is compar- 
atively rare, and occurs almost always in combination with 
oxygen and iron as chromic iron. This mineral, whose 
composition is represented by the formula FeCr 2 4 , may 
be regarded as the iron salt of an acid of the formula 
HCr0 2 . Replacing two atoms of hydrogen of this acid by 
one of iron, we should have a compound Fe(Cr0 2 ) 2 . This 
is analogous to spinel, which in a similar way is regarded 
as magnesium aluminate of the formula Mg(A10 2 ) 2 . 

Compounds of Chromium. — The principal compounds of 
chromium with which we have to deal are potassium chro- 
mate, K 2 Cr0 4 ; potassium dicliromate, K 2 Cr 2 7 , and other 
salts derived from chromic acid. There are, however, salts 
in which chromium takes the part of a metal, replacing the 
hydrogen of acids; as, for example, chromium sulphate, 
Cr 2 (S0 4 ) 3 . 

Potassium chromate, K 2 Cr0 4 , is formed when finely- 
powdered chromic iron is heated with potassium carbonate 
and potassium nitrate. 

Experiment 183. — Powder some chromic iron very fine. 
Thoroughly mix 3 grams with 3 grams each of potassium carbon- 
ate and potassium nitrate. Heat to fusion for some time in an 
iron crucible, using the blast-lamp. After cooling treat the mass 
with water, when a yellow-colored solution will be formed. Po- 
tassium chromate, K 2 CrC>4 , is in the solution. Save this solution. 

Potassium Dichromate, K 2 Cr 2 7 . — This is the form in 
which chromium is most frequently met with. It is formed 
from the chromate by adding acetic or nitric acid, and 



P0TIAS8UM BICHROMATE. 3TS 

forms large red crystals which are soluble in water. The 
change which takes place is represented thus: 

2K 2 Cr0 4 + 2HX0 3 = 2KX0 3 -I- £,0,0, + H 2 0. 

The relation between the chromate and the dichromate 
will be more readily understood by considering the acids 
from which they are derived. These are chromic acid, 
H 2 Cr0 4 , and dichromic acid, H 2 Cr 2 7 . The latter may be 
regarded as derived from the former by loss of water: 

2H 2 Cr0 4 = H 2 Cr 2 7 + H 2 0. 

The same relation exists between sulphuric acid, H 2 S0 4 , 
and disulphuric or fuming sulphuric acid, II,,S 2 7 . 

Experiment 184.— To the solution of potassium chromate 
already obtained add nitric acid enough to decompose the unused 
potassium carbonate and give the solution an acid reaction. The 
color will change from yellow to red. The red color indicates 
the presence of the dichromate. 

AVhen a solution of potassium dichromate is treated with 
potassium hydroxide until the color becomes pure yellow, 
the chromate is formed: 

K 2 Cr 2 7 + 2K0H = 2K 2 Cr0 4 + H 2 0. 

Experiment 185. — Convert 10 to 20 grams potassium dichro- 
mate into the chromate by the method mentioned. Evaporate to 
crystallization. 

The Chromate and Dichromate are good Oxidizing Agents. 
— Both the chromate and the dichromate are good oxidiz- 
ing agents. 

Experiment 186.— Treat a little of each salt in a test-tube with 
concentrated hydrochloric acid. [What evidence do you get that 
the salts are good oxidizing agents ?] 

(1) K 2 Cr04 + 8HC1 = 2KC1 + CrCl 3 +4 H 2 + 3C1. 

(2) K 2 Cr 2 7 4- 14HG1 = 2KCI + 2CrCl 3 + 7H 2 + 601. 



374 



INTRODUCTION TO CHEMISTRY. 



Insoluble Chromates. — The chromates of barium and 
lead, like the sulphates, are insoluble in water. They are 
yellow. The lead salt is the well-known pigment chrome- 
yellow. 

Experiment 187.— Add a little of a solution of potassium 
chromate or dichromate to a solution of a barium salt and of a 
lead salt. 



Chrome Alum is a salt allied to ordinary alum, but con- 
taining chromium instead of aluminium. Its formula is 
CrK(S0 4 ) 2 + 12H 2 0. The alums have analogous formulas. 



Ordinary alum A1K(S0 4 ), + 1211,0 

Iron alum FeK(S0 4 ), + 12H„0 

Chrome alum CrK(S0 4 ) a + 12H a 

Comparison of Chromium with Aluminium, Iron, and Sul- 
phur. — In its general chemical conduct chromium is similar 
to aluminium and iron on the one hand; while, on the 
other hand, its resemblance to sulphur is unmistakable, as 
is seen in the formation of the acids, chromic and dichro- 
mic acids, which are analogous to sulphuric and disulphuric 
acids, not only in composition, but in some of their prop- 
erties. L^re the lead and barium salts of sulphuric acid 
soluble or insoluble in water?] 

In its conduct towards reagents chromium more closely 
. resembles aluminium than iron. It forms no sulphide and 
no carbonate, so that when a soluble carbonate or sulphide 
is added to a solution of a chromium salt, such as chrome 
alum, the hydroxide is precipitated, as in the case of alu- 
minium. The hydroxide dissolves in caustic soda and 
caustic potash, but is reprecipitated when the solution is 
boiled. [How do aluminium and iron hydroxides act 
towards caustic soda ?] 

Experiment 188.— To a solution of potassium chromate add 
gome rather strong hydrochloric acid and a little alcohol. On 



CHEMICAL CONDUCT OF CHROMIUM— URANIUM. 375 

boiling the alcohol takes up oxygen from the chromate, a peculiai- 
smelling substance, aldehyde, is given off, and the solution now 
contains chromium chloride, CrCl 3 . The solution has a green 
color. The change is represented thus : 

2K 9 CrO, + 3C 2 HoO + 10HCI = 4KGM + 2CrCl,+ 3C 2 H 4 +8H 9 0. 

Alcohol. Aldehyde. 

To separate portions of the diluted solution add ammonium 
sulphide, sodium carbonate, and sodium hydroxide. The reac- 
tions which take place are: 

2CrCh + 3(NH0*S + 6H 2 = 2Cr(OH) a + 6XH 4 C1 + 3H 2 S. 
2CrCI 3 + 3Na a CO a 4- 3H 2 = 2Cr(OH) a + 6NaCI + 3CO a . 
CrCl 3 + 3NaOH = Cr(OH) 3 + 3NaCl. 

After noticing the general appearance of the precipitate formed 
with caustic soda, add an excess of the latter. [Does the precip- 
itate dissolve ?] Boil the solution. [What takes place ?] 

Uranium, U (At. Wt. 240). — This element occurs 
mostly in the form of the oxide U 3 8 known as pitch- 
blende. It forms salts in which the group UO a , called 
uranyl, takes the place of two atoms of hydrogen; as. for 
example, uranyl nitrate, U0 2 (X0 3 )., ; uranyl sulphate, 
U0 9 (S0 4 ). 

Uranium oxide, U,0 3 , conducts itself towards bases like 
iin acid, forming salts called uranates. 



CHAPTER XXVIII. 
- PALLADIUM.— PLATINUM.— GOLD. 

Palladium, ruthenium, and rhodium are three rare ele- 
ments which closely resemble one another. 

Palladium forms with hydrogen a compound which in 
general has the properties of alloys. It has the composi- 
tion Pd 2 H, and contains about GOO volumes of hydrogen to 
1 volume of palladium. The properties of this substance 
have led to the view that hydrogen has metallic properties. 
If by the name metal is meant an element that forms salts 
with acids, then it may be said that hydrogen bears to 
other metals a relation similar to that which carbonic acid 
bears to other acids. Acids are simply salts of hydrogen, 
and other metals drive out the hydrogen. Carbonates are 
in the same way decomposed by all other acids. 

Platinum, osmium, iridium, and gold form a group in 
which, however, the three first mentioned are the most 
closely related. Of these three, platinum is the best 
known. 

Platinum, Pt (At. Wt. 194.3), occurs almost always ac- 
companied by iridium, palladium, rhodium, ruthenium, 
and osmium, in the form of alloys. The ore is found in 
the Ural Mountains, in California, Australia, and a few 
other places. It is prepared by treating the ore with 
strong aqua regia, which dissolves the platinum, together 
with some iridium. The platinum chloride thus obtained 
is precipitated by means of ammonium chloride, with 

376 



PLATINUM. 377 

which, as with potassium chloride (see p. 299), it forms 
a difficultly soluble compound, PtCl 4 + 2NH 4 C1, or 
(NH 4 ) 2 PtCl 6 . When this is heated to a sufficiently high 
temperature it is decomposed, leaving metallic platinum 
as a residue. By special methods the iridium can be sepa- 
rated from it. 

Properties. — Platinum is a grayish-white metal, with a 
high lustre. Its specific gravity is 21.15, it being one of 
the heaviest substances known. The specific gravity of 
iron is 7.8, that of lead 11.4, and that of lithium 0.59. hi 
other words, a piece of platinum weighs nearly three times 
as much as a piece of iron of the same dimensions, and 
nearly twice as much as a piece of lead of the same dimen- 
sions. Platinum is not dissolved by hydrochloric, nitric, 
or sulphuric acid; but aqua regia dissolves it, forming 
platinum chloride, PtCl 4 . Fusing caustic alkalies attack 
it; sodium carbonate does not. It does not change in the 
air, and does not melt except in the flame of the oxy hydro- 
gen blowpipe. It resists the action of most substances. 
These properties make it extremely valuable to the chemist. 
Platinum crucibles and evaporating-dishes, foil, and wire 
are constantly used in the laboratory, and it is difficult to 
see how we could get along without them. Large retorts 
of platinum are used for the purpose of concentrating 
sulphuric acid and distilling it. 

Alloys of Platinum. — The only alloy of platinum of im- 
portance is that which it forms with iridium. A small 
percentage of this metal diminishes the malleability of 
platinum very markedly, and makes it brittle; but in- 
creases its power of resistance to the action of reagents. 
An alloy of 90 per cent platinum and 10 per cent iridium 
has been adopted by the French Government as the best 
material from which to make normal meters. This alloy 
is very hard, as elastic as steel, more difficultly fusible 



378 INTRODUCTION TO CHEMISTRY. 

than platinum, entirely unchangeable in the air, and is 
capable of a high polish. 

Platinum chloride, PtCl 4 , is made by dissolving the metal 
in aqua regia and evaporating off the acids. It dissolves in 
water, forming a yellowish-red solution, which is used in 
the laboratory for the purpose of precipitating potassium 
from its solutions, as the salt potassium chloroplatinate, 
K 2 PtCl fi or PtCl 4 + 2KC1, is difficultly soluble in water. 
The corresponding sodium salt, Na a PtCl 6 + 6II 2 0, is easily 
soluble in water. There is another chloride of platinum 
of the formula PtCl 2 , known as platinous chloride. 

Platinum Bases. — Platinum chloride combines with am- 
monia in a great many different proportions, forming the 
so-called platinum bases. The discussion of these com- 
pounds would lead us too far at present. 

Gold, An (At. Wt. 196.7). — According to the arrangement 
of the elements in the periodic system gold falls in the same 
group with copper and silver, to which it presents some 
points of resemblance. It forms more properly the con- 
necting link between the platinum group and the mem- 
bers of the second and third groups. 

Forms in which Gold occurs in Nature. — Gold is gener- 
ally found in nature in the native condition — a fact which 
is undoubtedly due to its chemical inactivity. That which 
is found in nature is never pure, but contains silver, and 
also, in different localities, iron, copper, and other metals. 
It is also found to some extent in combination with tellu- 
rium in the compounds AuTe 2 and (AuAg) 2 Te 3 . Native 
gold is frequently found enclosed in quartz, or more com- 
monly in quartz sand. The principal localities in which 
it is found are California and some of the other Western 
States, and Australia, Hungary, Siberia, and Africa. 



METALLURGY OF GOLD. 370 

Metallurgy of Gold. — From the chemical point of view 
the metallurgy of gold is in general very simple. There 
are two kinds of gold-mining — called placer-mining and 
vein-mining. In the former the earth and sand that con- 
tain gold are washed with water, which carries away the 
lighter particles, and leaves the gold mixed with other 
heavy materials. This mixture is then treated with mer- 
cury, which forms an amalgam with the gold, as it does 
with silver, and when this is placed in a retort and heated, 
the mercury passes over and leaves the gold behind. If 
silver is present, as is frequently the case, this is separated 
with the gold. In vein-mining the gold ores are taken out 
of veins in the earth, and the gold separated by grinding 
the ores and treating them with mercury, as in the last 
stage of placer-mining. Hydraulic mining is a modifica- 
tion of ordinary placer-mining. It consists in forcing 
water under pressure against the sides of hills and moun- 
tains in which gold occurs loosely mixed with the earth. 
The earth is thus carried away and the heavier gold is de- 
posited in sluices. 

The gold obtained as above is not pure. It can be 
separated from silver by dissolving it in aqua regia, evap- 
orating so as to drive off the nitric acid, then diluting, and 
treating with a reducing agent, when metallic gold is pre- 
cipitated. Thus when ferrous sulphate is used the follow- 
ing reaction takes place : 

3FeS0 4 + AuCl 3 = Fe,(S0 4 ) 3 + FeCl 3 + Au. 

Another method of separating silver from an alloy with 
gold consists in treating the metal with nitric acid or with 
boiling concentrated sulphuric acid, which dissolves the 
silver and leaves the gold. This process is not satisfac- 
tory, however, unless the amount of gold in the alloy is 
less than 25 per cent. If the proportion of gold is greater 
than this, the alloy is melted with silver enough to bring 



380 INTRODUCTION TO CHEMISTRY. 

the percentage of gold down to that mentioned. This is 
known as " quartation" 

Properties. — Gold is a yellow metal with a high lustre. 
It is quite soft and extremely malleable, so that it is possi- 
ble to make from it sheets the thickness of which is not 
more "than 0.000002 millimetre. Thin sheets are translu- 
cent, and the transmitted light appears green. Its spe- 
cific gravity is 19.3 ; its melting-point higher than that of 
copper, being about 1200°. It crystallizes in the regular 
system. Gold combines directly with chlorine, but not 
with oxygen. The three acids, hydrochloric, nitric, and 
sulphuric, do not act upon it ; but aqua regid dissolves 
it, forming auric chloride, AuCl 3 , in consequence of the 
evolution of nascent chlorine. Molten caustic alkalies and 
their nitrates act upon it, probably in consequence of the 
tendency to form aurates. 

Alloys of Gold. — The principal alloy of gold is that 
which contains copper. The standard gold coin of the 
United States contains nine parts of gold to one of copper. 
The composition of gold used for jewelry is usually stated 
in terms of carats. Pure gold is 24-carat gold ; 20-carat 
gold contains 20 parts of gold and 4 parts of copper; 18- 
carat gold contains 18 parts of gold and 6 parts of copper, 
etc. Copper gives gold a reddish color, and makes it harder 
and more easily fusible. Gold is also alloyed with silver ; 
and the alloy with mercury, known as gold-amalgam, is 
extensively used in the processes for extracting gold from 
its ores. 

Chlorides of Gold. — When gold is dissolved in aqua rer/ia 
ii is converted into auric chloride, AuCl 3 ; and if this 
solution is evaporated a part of the chloride is decom- 
posed into auroiis chloride, An CI, and chlorine. When 
gold is treated with dry chlorine it forms the dichloride, 



CHLORIDES OF GOLD. 381 

AuCl 2 . This, when treated with a little water, breaks 
down into auric chloride and aurous chloride, and by fur- 
ther treatment with water the latter yields auric chloride 
and gold. By filtering and evaporating to dryness, auric 
chloride is obtained in the anhydrous condition. It can 
also be obtained in crystallized form, the crystals having 
the composition AuCl 3 -f 2H 2 0. When anhydrous auric 
chloride is heated to 185°, it loses chlorine and is con- 
verted into aurous chloride, AuCl. This, as stated above, 
yields auric chloride and gold when treated with water. 
When treated with a solution of stannous chloride a solu- 
tion of auric chloride gives a purple-colored precipitate, 
known as the purple of Cassius, which appears to consist 
of finely-divided gold. 



I 



CHAPTER XXIX. 

SOME FAMILIAR COMPOUNDS OF CARBON. 

Organic Chemistry. — When the compounds that are ob- 
tained from plants and animals were first studied, they 
were supposed to be entirely different from the compounds 
obtained from the inorganic, or mineral, constituents of the 
earth. The former were called organic compounds because 
they were obtained from organized things; while the latter 
were called inorganic compounds. Organic compounds 
were the subject of Organic Chemistry, and inorganic com- 
pounds formed the subject of Inorganic Chemistry. These 
names are still in use, though they have lost their original 
meaning. Organic Chemistry now means only the Chem- 
istry of the Compounds of Carbon. 

Occurrence of the Compounds of Carbon. — The com- 
pound of carbon that occurs most widely distributed in 
nature is carbon dioxide. This, as has been pointed out, 
is the starting-point of all life on the globe. All living 
things are formed from it either directly or indirectly. 
Attention has been called to the fact that starch and 
cellulose are the principal compounds found in plants, and 
that fats, albumin, and fibrin are the most common sub- 
stances found in animals. 

Formation of Hydrocarbons. — Certain natural processes 
which are not thoroughly understood have given rise to the 

382 



DISTILLATION OF COAL. 383 

formation of a complex mixture of organic compounds, 
principally hydrocarbons, in petroleum. 

Distillation of Coal. — The destructive distillation of coal 
for the purpose of making illuminating-gas, and the forma- 
tion of coal-tar, have already been referred to. Coal-tar is 
one of the most important sources of compounds of carbon. 
The hydrocarbons benzene, C C H C , toluene, 0,11^ , xylene, 
CMI 10 , naphthalene, C 10 H a , anthracene, C 14 H ]0 ,ete., are 
obtained from this source. 

Distillation of Wood.— Wood is heated in closed vessels 
mostly for the purpose of making charcoal, as already ex- 
plained. Among the products obtained from this source 
are wood-spirit, or methyl alcohol, and pyrolignemts acid, 
or acetic acid. Large quantities of acetic acid are pre- 
pared in this way. 

Distillation of Bones. — In order to make bone-black, 
bones are subjected to destructive distillation. The oil 
which passes over is collected and known as bone-oil. This 
is the source of a large number of compounds which are of 
special interest on account of their connection with the 
valuable alkaloids quinine, morphine, etc. 

Fermentation.— A number of the most important com- 
pounds of carbon are formed by a process known as fer- 
mentation. This is a general term meaning any process in 
which a chemical change is effected by means of minute 
animal or vegetable organisms. The best-known example 
of fermentation is that of sugar, which gives rise to the 
formation of ordinary alcohol. 

Classes of Compounds of Carbon. — The chief classes of 
these compounds are the hydrocarbons; the alcohols; the 
aldehydes; the acids; the ethers; and the ethereal salts. 



384 INTRODUCTION TO CHEMISTRY. 

First a few of the best-known examples of each of these 
classes will be taken up, and afterwards some other familiar 
compounds which do not belong to any one of these 
classes. 



HYDROCARBONS. 

Compounds of Carbon and Hydrogen. — It is not an easy 
matter to effect combination between carbon and hydrogen 
in the laboratory except in a few simple cases. In nature 
processes are in operation which give rise to the formation 
of a large number of compounds containing these elements; 
and, further, in the manufacture of illuminating-gas from 
coal the conditions are such as to cause the combination of 
carbon and hydrogen, several interesting compounds being 
thus formed. There are no other two elements which 
combine with each other in as many different proportions 
as carbon and hydrogen. The compounds thus formed are 
known as hydrocarbons. The number of hydrocarbons 
known is very great, being somewhere near two hundred. 
Fortunately, investigation has shown that quite simple re- 
lations exist between these compounds; and hence, though 
the number is large, the study is not as difficult as might 
be expected. 

Petroleum is an oily liquid found in many places in the 
earth in large quantity, particularly in Pennsylvania and 
the Caucasus, In the earth it contains both gases and liq- 
uids. When it is brought into the air, the presssure being 
removed, the gases are given off. There are several gas- 
eous hydrocarbons given off, and a large number of liquids 
left behind. 

Refining of Petroleum. — The vapors from petroleum 
when mix.ed with air are explosive, and the thicker liquids 
clog the lamps and wicks. Therefore these must be re- 



PETROLEUM-IIOMOLOG Y. 3S5 

moved before the oil is fit for household use. This is done 
by (1) distilling, (2) washing with sulphuric acid, (3) wash- 
ing with alkali, and (4) washing with water. The product 
thus prepared is called kerosene. 

In refining petroleum a number of products are obtained 
which cannot be used in lamps. Those which are lighter 
than kerosene, that is to say those which boil at a lower 
temperature, are known as gasoline, naphtha, benzine, etc. 
From the heavier portions, or those which boil at higher 
temperatures than kerosene, paraffin is made. Each of 
these substances is a mixture of several chemical com- 
pounds. 

Hydrocarbons contained in Petroleum. — The simplest 
hydrocarbon contained in petroleum is methane, or marsh- 
gas, CH 4 ; the next has the composition CJI 6 , the next 
C 3 H H ,etc. It will be seen that these compounds bear a 
simple relation to one another, as far as composition is 
concerned. They are the first members of a series the 
names and symbols of the first eight members of which are 
given below : 



e 



CH 4 , 


Methane, or Marsh-gas; 




C 2 H,, 


Ethane; 


i 


C 3 H e , 


Propane ; 


1 


C 4 H,„, 


Butane; 




C a H 12 , 


Pentane ; 




OH,, 


Hexane ; 




0,H ltJ 


Heptane; 




c e n iS , 


Octane. 





Homology. — The first member of the series differs from 
the second by CH 2 ; there is also this same difference, in 
general, between any two consecutive members of the 
series. This relation is known as homology, and such a 
series as an homologous series. Carbon is distinguished 



386 



INTLiObUCTlON TO CllEMISTRr. 



from all other elements by its power to form homologous 
series. 

The Ethylene Series of Hydrocarbons.— Besides the series 
above mentioned, which is known as the marsh-gas series, 
there are other homologous series of hydrocarbons. There 
is one beginning with ethylene, C 2 H 4 , examples of which 
are 

Ethylene, CJI 4 ; 
Propylene, C 3 H 6 ; 
Butylene, C 4 H B . 

The Acetylene Series.— There is a series beginning with 
acetylene, examples of which are 

Acetylene, C 2 H 2 ; 
Allylene, 3 H 4 . 

The Benzene Series. — Another series begins with benzene, 
C ( .H R . Some of the members of this series are 



Benzene, C 6 H e ; 
Toluene, C 7 H 8 ; 
Xylene, C a H 10 . 



Marsh-gas, Methane, Fire-damp, CH 4 . — Marsh-gas is 
found in nature in petroleum, and is given off when the 
oil is taken out of the earth, and the pressure is removed. 
It is formed, as the name implies, in marshes, as the prod- 
uct of a reducing process. Vegetable matter is composed 
of carbon, hydrogen, and oxygen. When it undergoes de- 
composition in the air in a free supply of oxygen, the final 
products formed are carbon dioxide and water. When the 
decomposition takes place without access of oxygen, as 
under water, marsh-gas, which is a reduction-product, is 
formed. The gas can be made in the laboratory by pass- 
ing a mixture of hydrogen sulphide, H 2 S, and the vapor of 






MARSH GAS-SUBSTITUTION PROD UCT& 387 

carbon bisulphide, CS 2 , over heated copper. The sulphur 
is extracted from the compounds, and the carbon and hy- 
drogen combine, as represented in the equation 

CS 2 + 2HJS + SCu = CH 4 + 4Cn,S. 

Marsh-gas is met with in coal-mines, and is known to 
the miners as fire-damp, damp being the general name 
applied to a gas, and the name lire-damp meaning a gas 
that burns. To prepare it in the laboratory, it is most 
convenient to heat a mixture of sodium acetate and quick- 
lime. The change which takes place will be most readily 
understood by considering it as a simple decomposition of 
acetic acid. Acetic acid has the formula C,H 4 0,. When 
heated alone, it boils and does not suffer decomposition. 
If it is converted into a salt, and heated in the presence of 
a base, it breaks up into marsh-gas and carbon dioxide: 

C.JI 4 2 = CH 4 + CO,. 

The carbon dioxide, which forms salts with bases, does not 
pass off, but remains behind in the form of a salt of car- 
bonic acid. 

Marsh-gas is a colorless, transparent, tasteless, inodorous 
gas. It is slightly soluble in water. It burns, forming 
carbon dioxide and water. AVhen mixed with air, the 
mixture explodes if a flame or spark comes in contact with 
it. This is one of the causes of the explosions which so 
frequently occur in coal-mines. To prevent these explo- 
sions a special lamp was invented by Sir Humphry Davy, 
which is known as Davy's safety-lam}) (p. 1S3). 

Substitution-products of the Hydrocarbons. — Marsh-gas 
and other hydrocarbons undergo change when treated with 
chlorine and bromine. The change consists in the substitu- 
tion of one or more atoms of chlorine or of bromine for the 



388 INTRODUCTION TO CHEMISTRY. 

same number of atoms of hydrogen. In the case of marsh- 
gas and chlorine the possible changes are represented as 
below : 

CH 4 + 01, = CH 3 C1 + HC1; 

OH3OI + 01, = CH 2 C1 2 + HOI; 

CH 2 C1 2 + 01, = CHOI3 + HOI; 

CHCI3 + 01, = CC1 4 + HOL 

All the products represented are known. 

Chloroform, CHC1 3 . — Chloroform can be made as above 
indicated, but it is made on the large scale by treating 
alcohol (which see) or acetone (which see) with bleaching- 
powder. It is a heavy liquid with an ethereal odor and a 
somewhat sweet taste. It is one of the most valuable 
anaesthetics, though there is some danger attending its use. 

Iodoform, CHI 3 . — This compound, like chloroform, is a 
substitution-product of marsh-gas. It is made by bring- 
ing together alcohol, an alkali, and iodine. It is a solid 
substance, soluble in alcohol and ether, but insoluble in 
water. It crystallizes in six-sided yellow plates. It is ex- 
tensively used as a dressing for wounds in surgery. 

Ethylene, Olefiant Gas, 0,H 4 . — This hydrocarbon is 
formed by heating a mixture of ordinary alcohol and con- 
centrated sulphuric acid. The reaction is represented 
thus: 

2 H 6 = H,0 + 2 H 4 . 

Alcohol. Ethylene. 

r i> 

Ethylene is a colorless gas, which can be condensed to a 
liquid. It burns with a luminous flame. With oxygen it 
forms an explosive mixture. 

Acetylene, C 2 H 2 . — Acetylene is formed when a current 
of hydrogen is passed between carbon poles, which are in- 
candescent in consequence of the passage of a powerful 



METHYL ALCOHOL-ETHYL ALCOHOL. 



389 



electric current. In this case carbon and hydrogen com- 
bine directly. It is formed also when the flame of an 
ordinary laboratory gas-burner (Bunsen burner) " strikes 
back/' or burns at the base without a free supply of air. 
Its odor is unpleasant. It burns with a luminous, smoky 
flame. 

ALCOHOLS. 

Methyl Alcohol, Wood-spirit, CH 4 0. — This is formed in 
the distillation of wood, and must be separated from the 
other products which are formed at the same time. It has, 
when pure, a pleasant odor and taste, and acts upon the 
animal system very much as ordinary alcohol does. It burns 
without giving light or smoke, and may therefore be used 
in lamps for heating-purposes as ordinary alcohol is. It is 
used in the manufacture of varnishes. 



Ethyl Alcohol, Spirits of Wine, C a H O— This well- 
known substance is formed by the fermentation of grape- 
sugar or glucose. 

Experiment 188. — Dissolve about 150 grams of commercial 
.©•rape-sugar in 1| litres of water in a flask. Connect the flask by 
means of a bent glass tube 
with a cylinder or bottle con- 
taining clear lime water. The 
vessel containing the lime- 
water must be provided with a 
cork with two holes. Through 
one of these passes the tube 
from the fermentation-flask ; 
through the other a tube con- 
necting with a vessel contain- 
ing solid caustic potash, the object of which is to prevent the air 
from acting upon the lime-water. The arrangement of the appa- 
ratus is shown in Fig. 56. Now add to the solution of grape- 
sugar a little fresh brewer's veast ; close the connections and 




Fig. 5G. 



390 INTRODUCTION TO CHEMISTRY. 

and allow to stand. 80011 an evolution of gas will begin, and, as 
this passes through the lime-water, a precipitate will be formed 
which can be shown to be calcium carbonate. 

What Change takes Place in the Sugar? — If the solution 
in the flask is examined carefully it will be found to con- 
tain alcohol and no sugar. Grape-sugar has the composi- 
tion expressed by the formula C 6 H I2 O fi . By fermentation 
it is decomposed, forming alcohol, C 2 1I 6 0, and carbon 
dioxide, C0 2 , thus: 

C 6 H J2 6 = 2C 2 H 6 + 2C0 2 . 

What Causes the Change? — It has been found that the 
change of grape-sugar is caused by small organized bodies 
which grow in the solution. These bodies are contained in 
ordinary yeast. 

Germs in the Air. — When fruit-juices which contain 
sugar are exposed to the air they undergo fermentation 
without the addition of yeast. This is due to the fact that 
the germs or seeds of the bodies which cause fermentation 
are everywhere floating in the air. Hence when a liquid 
in which these seeds can grow is exposed to the air, the 
bodies are formed and fermentation takes place. 

Different Kinds of Fermentation. — The fermentation 
which yields alcohol is only one of many kinds. Among 
the others are: (1) lactic-acid fermentation, which takes 
place in the souring of milk; and (2) acetic-acid fermenta- 
tion, which causes the transformation of alcohol into acetic 
acid. The. latter ferment is contained in "mother of 
vinegar." 

Distillation of Fermented Liquids. — In order to get the 
alcohol from liquids which have undergone fermentation 
they must be distilled. For this purpose very perfect 
forms of stills have been devised, so that the alcohol passes 



ALCOHOL - G L YCER1N. 39 1 

over nearly free from other substances. Usually it con- 
tains impurities known as fusel oil. 

Properties of Alcohol. — Pure ethyl alcohol has a peculiar, 
pleasant odor. It remains liquid at very low temperatures, 
but has been converted into a solid at a temperature of 
— 130.5°. It burns with a flame which does not deposit 
soot, and was hence formerly much used in laboratories 
for heating purposes, and is still used where gas cannot 
be obtained. Its effects upon the human system are Avell 
known. It intoxicates when taken in dilute form, while 
in large doses it is poisonous. It lowers the temperature 
of the body when taken internally, although it causes a 
sensation of warmth. 

Uses of Alcohol. — Alcohol is the principal solvent for or- 
ganic substances. It is hence extensively used in the arts, 
as in the manufacture of varnishes, perfumes, and tinctures 
of drugs. Most beverages in use owe their intoxicating- 
power to the presence of alcohol. The milder forms of 
beer contain from 2 to -3 per cent; light wines about 8 
per cent; while whiskey, brandy, etc., sometimes contaii: 
as much as 60 to 75 per cent. 

Glycerin, C 3 H H 3 . — Glycerin is an alcohol which occurs 
very widely distributed as a constituent of fats. The rela- 
tion it ber.rs to the fats will be explained when the acids 
which enter into the fats are taken up. It is obtained from 
the fats by boiling them with an alkali like caustic soda or 
caustic potash, or by heating with steam. 

Properties. — Glycerin is a thick, colorless liquid with a 
sweetish taste. It attracts moisture from the air, and is 
hence used to keep surfaces moist. 



392 INTRODUCTION TO CHEMISTRY. 



ALDEHYDES. 

Acetic Aldehyde, Ordinary Aldehyde, C 2 H 4 0. — This com- 
pound is formed by oxidizing ordinary alcohol, the change 
being represented by this equation: 

C 2 H 6 + = C 2 H 4 + H 2 0. 

Experiment 189. — In a small flask put a few pieces of potassium 
dichromate, K 2 Cr 2 7 , and pour upon it a few cubic centimetres of 
moderately concentrated sulphuric acid. To this mixture add 
slowly a few cubic centimetres of ordinary alcohol. Notice the 
odor. 

Aldehyde is a volatile liquid with a characteristic pene- 
trating odor. When left to itself, and especially when 
treated with a number of other things, it is converted into 
another substance of the same composition. This is called 
paraldehyde. A determination of the molecular weight of 
the substance by the method of Avogadro has shown that it 
must be represented by the formula C 6 H 12 3 . The change 
from aldehyde to paraldehyde must, therefore, be repre- 
sented thus: 

3C 2 H 4 = C 6 H 12 3 . 

Paraldehyde is used in medicine. 

Chloral, C 2 Cl 3 HO, is a compound formed by the action 
of chlorine on alcohol. It is related to aldehyde, as chloro- 
form is related to marsh-gas, that is to say, it is a trichlorine 
substitution-product. It is a colorless liquid. With water 
it forms a crystallized compound, chloral hydrate, C 2 Cl,HO 
+ H 2 0, which is easily soluble in water, and crystallizes 
from the solution in colorless prisms. Taken internally in 
doses of from 1.5 to 5 grams, it produces sleep. In larger 
doses it acts as an anaesthetic, 



FORMIC ACID-ACETIC ACID. 393 



ACIDS. 

Formic Acid, CH 2 2 . — This acid occurs in nature in red 
aufcg, in stinging-nettles, in the shoots of some of the vari- 
eties of pine, and elsewhere. It is a colorless liquid. 
Dropped on the skin, it causes extreme pain and produces 
blisters. 

Acetic Acid, C,H 4 2 . — This is the acid contained in vine- 
gar, and the value of vinegar is due to its presence. It is 
formed from alcoholic liquids by exposing them to the air, 
in consequence of the presence of a microscopic organism 
which is contained in w T hat is commonly known as "mother 
of vinegar." The formation of acetic acid from alcohol is 
due to the action of oxygen as represented in the equation 

C 2 H 8 + 2 = 2 H 4 2 + H 2 0. 

Alcohol. Acetic acid. 

But oxygen alone does not effect the change. When the 

ferment is present the oxidation takes place. Acetic acid 

is also obtained by distilling wood. Hence the names 
pyroligneous acid and wood-vinegar. 

Properties. — Acetic acid is a clear, colorless liquid. It 
has a very penetrating, pleasant, acid odor, and a sharp 
taste. The pure substance acts upon the skin like formic 
acid, causing pain and raising blisters. 

Uses. — Acetic acid is extensively used, chiefly in the dilute 
form known as vinegar. It is used in calico-printing in 
the form of iron and aluminium salts. With iron it gives 
hydrogen, which is needed in the manufacture of certain 
compounds used in making dyes. 

Salts of Acetic Acid. — The best-known salts of acetic 
acid are lead acetate, Pb(C 2 H 3 2 ) 2 , commonly called sugar 
of lead; and copper acetate, Cu(C 2 H 3 2 ) 2 , a variety of 
which is known as verdigris. 



894 INTRODUCTION TO CHEMISTRY. 

Fatty Acids. — Formic and acetic acids are the first mem- 
bers of an homologous series (see page 385). Some of the 
more important members are named in the following table : 

Formic acid CH 2 2 . 

Acetic " C 2 H 4 2 . 

Propionic " C 3 H 6 O a . 

Butyric " ... C 4 H 8 2 . 

Palmitic " C 16 H 32 02. 

Stearic " .. Ci 8 H 36 2 . 

They are called fatty acids for the reason that many of 
them are obtained from fats. 

Butyric acid, C 4 H 8 2 , is of special interest because it is 
obtained from butter by boiling with caustic potash. It 
occurs also in many other fats. There is a butyric-acid 
ferment contained in putrid cheese which has the power of 
converting sugar into butyric acid. 

Palmitic acid, C 16 H 32 2 , is obtained from many fats, 
but palm-oil is especially rich in it. 

Stearic acid, C 18 H 36 2 , is the acid contained in the fat 
known as stearin. The so-called "stearin candles" are 
made of a mixture of palmitic and stearic acids. 

Soaps. — Soaps are the alkali salts of the acids contained 
in fats, especially of palmitic and stearic acids. Fats are 
compounds of these acids with glycerin. When the fats 
are boiled with an alkali, as caustic soda, the corresponding 
salts of the acids are formed, while the glycerin is set free. 
The palmitate and stearate of potassium and sodium are 
the soaps. 

Experiment 190. — In an iron pot boil a quarter of a pound of 
lard with a solution of 40 grams caustic soda in 250 cc. of water 
for an hour or two. After cooling add a strong solution of sodium 
chloride. The soap formed will separate and rise to the top of 
the solution, where it will finally solidify. Dissolve some of the 
soap thus obtained in water. 



SOAP. 395 

TTse of Soap. — The cleansing power of soap depends upon 
the fact that it dissolves the oily film on the surface of the 
skin and thus facilitates the removal of the foreign sub- 
stances commonly known as dirt. 

Action of Soap on Hard Waters. — As has been explained, 
a hard water is one that contains salts in solution. Tem- 
porary hardness is that which is caused by calcium car- 
bonate held in solution in the water by carbon dioxide. 
Permanent hardness is caused by calcium sulphate or mag- 
nesium salts. The calcium and magnesium salts of pal- 
mitic and stearic acids are insoluble in water. Therefore, 
when soap is added to a hard water these insoluble salts 
are precipitated and give the water a hard feeling. In 
attempting to wash the hands with soap in a hard water 
they become covered with a thin layer of the insoluble salts 
which prevents them from rubbing freely over each other, 
and makes them feel sticky. Before the soap can do any 
good all the lime-salt must be precipitated. The action in 
the case of temporary hardness is represented by the equa- 
tion 

2NaC 16 H 3) 0, + CaC0 3 = Ca(C 1 .H, 1 1 ) 1 + Na,C0 3 . 

Soap. Calcium palmitate. 

In the case of permanent hardness it is represented by 
the equation 

2NaC 16 H s ,0 2 + CaS0 4 = Ca(C 16 H 3l 2 ) 2 + Na 2 SO,. 

Experiment 191. — Make some hard water by passing carbon 
dioxide through dilute lime-water until the precipitate first formed 
is dissolved again. Filter. Make a solution of soap by shaking 
up a few shavings of soap with water. Filter. Add the solution 
of soap to the hard water. Is a precipitate formed ? Rub a piece 
of soap between the hands wet with the hard water. 

Experiment 192. — Make some hard water by shaking a litre or 
two of water with a little powdered gypsum. Perform with it the 



396 INTRODUCTION TO CHEMISTRY. 

same experiments as those first performed with the water con- 
taining calcium carbonate. 

Relations of the Soap Industry to other Industries. — A 

great chemist and philosopher has said that the quantity o£ 
soap used in a country is a measure of the civilization of 
that country. Certain it is that soap is only used by 
civilized people, and that by them it is used in enormous 
quantities. In many farm-houses a primitive method for 
the manufacture of soap is practised, consisting in treating 
refuse fats with the lye extracted from wood-ashes. A soft 
soapy mass is thus obtained known as "soft-soap." Fats 
form the starting-point in the manufacture of all soap. 
These are generally treated with caustic soda. Caustic 
soda is all made from sodium carbonate by the action of 
lime; and, as has been seen, sodium carbonate is made 
from common salt mostly by the Leblanc process, which 
requires sulphuric acid. Thus the manufacture of sul- 
phuric acid and sodium carbonate is intimately related to 
the manufacture of soap. 

Oxalic Acid, C 2 H 2 4 . — This acid occurs very widely dis- 
tributed in nature, as in the sorrels, which owe their acid 
taste to the presence of acid potassium oxalate, KC 2 H0 4 ; 
and as the ammonium salt in guano. It is probably one of 
the first substances formed from carbon dioxide in the plant. 
It is manufactured by heating wood shavings or sawdust 
with caustic soda and caustic potash. Oxalic acid is an 
active poison. It is used in calico-printing, and in cleaning 
brass and copper surfaces. 

Lactic Acid, C 3 H 6 3 . — Lactic acid is made by the fer- 
mentation of sugar by means of the lactic-acid ferment. 
The reaction effected by the ferment is represented by the 
equation 

C 6 H ]2 6 = 2C 3 H 6 0, 



MALIC, TARTARIC, AND CITRIC ACIDS— ETHER. 397 

Malic Acid, C 4 H 6 6 . — This acid is very widely distributed 
in the vegetable kingdom, as in apples, cherries, etc. 

Tartaric Acid, C 4 H 6 6 . — Tartaric acid occurs very widely 
distributed in fruits, sometimes uncombined, sometimes in 
the form of the potassium or calcium salt; as, for example, 
in grapes, berries of the mountain-ash, potatoes, cucum- 
bers, etc., etc. It is prepared from " cream of tartar." 
This is acid potassium tartrate, which is formed when 
grape-juice ferments. 

Citric Acid, C 6 H B 7 . — Citric acid, like malic and tartaric 
acids, is very widely distributed in nature in many varieties 
of fruit, especially in lemons. It is also found in currants, 
whortleberries, raspberries, gooseberries, etc., etc. It is 
prepared from lemon-juice : 100 parts of lemons yield 5£ 
parts of the acid. It is a solid, crystallized substance, sol- 
uble in water. It is frequently used for the purpose of 
making lemonade without lemons, and there is no objection 
to its use for this purpose. 

ETHERS. 

Ether, C 4 H 10 O. — Ordinary ether is the best-know r n repre- 
sentative of the class of compounds called ethers. It is 
formed from ordinary alcohol by treating it with sulphuric 
acid and distilling. The result of the action which takes 
place is represented by the equation 

2C 2 H 6 = C 4 H 10 O + H,0. 

Alcohol. Ether. 

Ether is a liquid which boils at a low temperature and 
takes fire and burns readily. Inhaled it produces insensi- 
bility to pain. It is therefore called an ancesthetic. 

ETHEREAL SALTS. 

Action of Acids upon Alcohols. — When an acid acts upon 
an alcohol it is neutralized, though not as readily as when 



398 INTRODUCTION TO CHEMISTRY. 

it acts upon a base. The product is a substance which re- 
sembles a salt and is called an ethereal salt. Thus when 
nitric acid acts upon alcohol this reaction takes place : 

C 2 H 6 + HN0 3 = C 2 H s N0 3 + H 2 0. 

The product C 2 H 5 N0 3 , called ethyl nitrate, is an ethereal 
salt. The alcohol acts as if it were a substance like caus- 
tic potash and made up thus, C 2 H 5 OH. The resemblance 
between its action and that of caustic potash is shown by 
the equations 

KOH + HN0 3 = KN0 3 + H 2 0, and 
C,H 5 OH + HN0 3 - C.H.NO, + H 2 0. 

Saponification. — When an ethereal salt is boiled with a 
caustic alkali it is decomposed, the products being an alco- 
hol and an alkali salt. Thus when ethyl nitrate is boiled 
with caustic potash, potassium nitrate and alcohol are 
formed : 

C 2 H 5 N0 3 + KOH = C 2 H 5 OH + KN0 3 . 

This process is called saponification, because the most 
important example is furnished by soap-making. 

Fats. — The fats are ethereal salts in the formation of which 
glycerin, as the alcohol, and three acids take part. The 
three acids are palmitic and stearic acids, already mentioned, 
and oleic acid. C 18 H 34 2 . Although the composition of 
these substances is comparatively complex, the way they act 
upon one another is simple, and is the same as the action 
of nitric acid upon alcohol in forming ethyl nitrate. The 
fats, then, are the palmitate, stearate, and oleate of glyceryl, 
which bears to glycerin very much the same relation that 
ethyl, C 2 H 5 , bears to alcohol. When a fat is boiled with 
caustic soda, glycerin and the sodium salts of the acids con- 
tained in the fat are formed. 



ETHEREAL SALTS. 399 

Butter consists of ethereal salts of glycerin and several 
fatty acids, among which are palmitic, stearic, and butyric. 
Oleomargarin is an artificial butter made from other fats 
than that of milk. 

Ethereal Salts as Essences. — The ethereal salts generally 
have pleasant odors, and it is to their presence that many 
fruits owe their flavors. Some of the compounds are now 
made artificially and used instead of the fruit-extracts. 
Thus the ethyl salt of butyric acid is used under the name 
of essence of pineapples, and the amyl salt of valeric acid 
under the name of essence of apples. 

Nitroglycerin. — Among the more important ethereal 
salts of glycerin are the nitrates. Two of these are known, 

( O.N0 2 
viz., the mono-nitrate, C 3 H 5 \ OH , and the tri-nitrate, 

(OH 
C 3 H 5 (O.N0 2 ) 3 , the latter being the chief constituent of 
nitroglycerin. Nitroglycerin is prepared by treating gly- 
cerin with a mixture of concentrated sulphuric and nitric 
acids. It is a pale yellow oil which is insoluble in water. 
At — 20° it crystallizes in needles. It explodes very vio- 
lently by concussion. It may be burned in an open vessel, 
but if heated above 250° it explodes. Dynamite is infu- 
sorial earth * impregnated with nitroglycerin. Nitroglyc- 
erin is the active constituent of a number of explosives. 

RELATIONS BETWEEN" THE COMPOUNDS CONSIDERED. 

Comparison of the Formulas. — On comparing the formu- 
las of the hydrocarbons of the marsh-gas series (see page 
385) with those of the simplest alcohols and the fatty acids, 

* That is to say, earth made up of the microscopic flinty shells 
which constitute the fossil remains of certain minute and simple 
plants. 



400 INTRODUCTION TO CHEMISTRY. 






it will be seen that these compounds are all related in a 
simple way. Below are lists of a few of the hydrocarbons, 
alcohols, and acids: 



Hydrocarbons. 


Alcohols. 


Acids. 


CH 4 


CH 4 


CH 2 2 


C 2 H 6 


C 2 H 6 


C 2 H 4 0, 


8 H 8 


C 3 H 8 


C,H 6 0, 


C 4 H 10 ,etc. 


C 4 H 10 O, etc. 


C 4 H 8 2 , etc 



Each of these series is an homologous series. 

Alcohols. — Alcohols have been shown to be derived from 
the hydrocarbons by the replacement of one or more hydro- 
gen atoms by oxygen and hydrogen, OH, or from water 
by replacing one of the hydrogen atoms of the water by a 
compound of carbon and hydrogen. An alcohol, then, is 
a hydroxide, just as a metallic base is; only, instead of con- 
sisting of a metal in combination with hydrogen and oxy- 
gen, it consists of a compound of carbon and hydrogen in 
combination with hydrogen and oxygen. Thus: 

Metallic Bases. Alcohols. 

K(OH) CH 3 (OH) 

Na(OH) C 2 H 5 (OH) 

More Complex Alcohols. — Just as lime is a more complex 
base than caustic potash, as shown by the formulas KOH 
and Ca0 2 H 2 or Ca(OH),, so there are more complex alco- 
hols than ordinary alcohol. A good example is furnished 
by glycerin, C 3 H 8 3 , which has been shown to be a hydrox- 
ide corresponding to aluminium hydroxide, Al(OH) 3 , a 
fact which is represented by the formula C 3 H 5 (OH) 3 . It 
may be called glyceryl hydroxide, the complex, C 3 H., 
being known as glyceryl. 

Radicals or Residues. — The compounds of hydrogen and 
carbon contained in the alcohols are called radicals or rest- 



FORMULAS OF ACIDS. 401 

dues, We may say that an alcohol is water in which half 
of the hydrogen has been displaced by a radical. 



HOH 


C 2 H 5 OH 


Water. 


Ordinarv alcohol. 


HOH 


(OH 
C,H B ] OH = C.H.O, 

( OH 


HOH 


HOH 


Water. 


Glycerin. 



Acids. — Just as the alcohols have been shown to be de- 
rived from water, so the organic aoids have been shown to 
be derived from carbonic acid. Carbonic acid itself is not 
known. But the carbonates are derived from an acid of 

\ OTT 
the formula H 2 C0 3 , or CO - qtt. If, in this acid, a hy- 
droxy! is replaced by a radical, as, for example, by ethyl, 

c n TT 

C 2 H 5 , a substance of the formula CO ■] qV 5 or C 3 H 6 2 is 

the result. If methyl, CH 3 , is introduced in place of 

( PIT 
ethyl, the product is CO •] qtt 3 or C 2 H 4 o , which is acetic 

acid. In a similar way all the organic acids are derived 
from carbonic acid. 



CHAPTER XXX. 
OTHER COMPOUNDS OF CARBON. 

The Carbohydrates. — The carbohydrates form an im- 
portant group of carbon compounds which include the 
most abundant substances found in the vegetable kingdom. 
Besides carbon, they contain hydrogen and oxygen in the 
proportions to form water. Hence they are called carbo- 
hydrates. The chief compounds included under this head 
are grape-sicgar or glucose, cane-sugar, starch, cellulose, 
gum, and dextrin. 

Grape-sugar, Glucose, Dextrose, C 6 H 12 6 .— Dextrose oc- 
curs very widely distributed in the vegetable kingdom, 
particularly in sweet fruits. It is found also in honey and, 
further, in the liver and the blood. 

Formation of Dextrose. — Dextrose or glucose is formed 
from several of the carbohydrates by boiling with dilute 
mineral acids, or by the action of ferments. Its forma- 
tion from cane-sugar takes place according to this equa- 
tion, equal quantities of dextrose and levulose being 
formed : 

O it H„O n + H 2 = C 6 H 12 6 + C.H.,0.. 

Cane-sugar. Dextrose. Levulose. 

Its formation from starch is represented by this equa- 
tion; 

C 6 H 10 O 6 + H 2 = C 6 H 12 6 . 

Starch. Dextrose. 

403 



DEXTROSE-LE VULOSE- CANES UGAR. 403 

Manufacture of Dextrose or Glucose. — Dextrose is pre- 
pared on the large scale from corn-starch in the United 
States, and from potato-starch in Germany. The change 
is usually effected by boiling with dilute sulphuric acid. 
The acid is afterwards removed by treating with chalk, and 
filtering. [Explain how this removes the acid.] The 
filtered solutions are evaporated either to a syrupy consist- 
ency, and sent into the market under the names "glu 
cose," " mixing-syrup/' etc.; or to dryness, the solid prod- 
uct being known as " grape-sugar." 

Properties. — Dextrose crystallizes from concentrated so- 
lutions, and as seen in commercial " granulated grape- 
sugar" looks very much like granulated cane-sugar. It 
is sweet, but not as sweet as cane-sugar. It is estimated 
that the sweetness of dextrose is to that of cane-sugar as 
3 : 5. Under the influence of yeast it ferments, yielding 
mainly alcohol and carbon dioxide. Putrid cheese trans- 
forms it into lactic acid, and then into butyric acid. 

Levulose, Fruit-sugar, C 6 H 12 6 . — This form of sugar 
occurs with dextrose in fruits; and is formed by the action 
of dilute acids or ferments on cane-sugar, which breaks up 
according to the equation 

C lt H M O n + H 2 = C 6 H 12 6 + C.H„0.. 

Cane-sugar. Dextrose. Levulose. 

As cane-sugar is found in unripe fruits, it is probable 
that the change represented in the equation takes place 
during the process of ripening. 

Cane-sugar, C ]2 H 22 O n . — This well-known variety of sugar 
occurs very widely distributed in nature — in sugar-cane, 
sorghum, the Java palm, the sugar-maple, beets, madder- 
root, coffee, walnuts, hazel-nuts, sweet and bitter almonds; 
in the blossoms of many plants, etc., etc. 



404 INTRODUCTION TO CHEMISTRY. 

Sugar-refining. — Sugar is obtained mainly from the 
sugar-cane and beets. In either case the processes of ex- 
traction and refining are largely mechanical. When sugar- 
cane is used, this is macerated with water to dissolve the 
sugar. Thus a dark-colored solution is obtained. This is 
evaporated, and then passed through filters of bone-black 
by which the color is removed. The clear solution is then 
evaporated in open vessels to some extent; and, finally, in 
large closed vessels called " vacuum-pans," from which the 
air is partly exhausted, so that the boiling takes place at a 
lower temperature than is required under the ordinary 
pressure of the atmosphere. The mixture of crystals and 
mother-liquors obtained from the " vacuum-pans " is freed 
from the liquid by being brought into the " centrifugals." 
These are funnel-shaped sieves which are revolved rapidly, 
the liquid being thus thrown by centrifugal force through 
the openings of the sieve, while the crystals remain behind 
and are thus nearly dried. The final drying is effected by 
placing the crystals in a warm room. 

Molasses. — The mother-liquors obtained from the " cen- 
trifugals " are further evaporated, and yield lower grades 
of sugar; and, finally, a syrup is obtained which does not 
crystallize. This is molasses. 

Properties of Sugar. — Sugar crystallizes from water in 
large well-formed prisms. When heated to 210° to 220°, 
it loses water, and is converted into a substance called 
caramel, which is colored more or less brown. AVhen 
boiled with dilute acids, cane-sugar is split into equal parts 
of dextrose and levulose. The mixture of the two is called 
invert-sugar. Yeast gradually transforms cane-sugar into 
dextrose and levulose, and these then undergo fermenta- 
tion. Cane-sugar does not ferment. 

Sugar of Milk, Lactose, C )2 H 22 O n + H 2 0. — This sugar 
occurs in the milk of all mammals. It is obtained in the 



LACTOSE-CELL ULOSE. 405 

manufacture of cheese. Cow's milk consists of water, 
casein, butter, sugar of milk, and a little inorganic ma- 
terial, in about the following proportions: 

Water 87 per cent. 

Casei'11 4 

Butter 3£ 

Sugar of milk 4| ' ' 

Mineral matter f " 

Too 

Cheese is made by adding rennet to milk, which causes 
the separation of the casein. The sugar of milk remains 
in solution, is separated by evaporation, and purified by 
recrystallization. It has a slightly sweet taste, and is much 
less soluble in water than cane-sugar. 

Souring of Milk. — Sugar of milk ferments under certain 
circumstances, and is transformed mostly into lactic acid. 
The souring of milk is a result of this fermentation. The 
lactic acid formed coagulates the casein; hence the thick- 
ening. 

Cellulose, C 6 H 10 O 6 . — Cellulose forms, as it were, the 
groundwork of all vegetable tissues. It presents different 
appearances and different properties, according to the 
source from which it is obtained; but these differences are 
due to substances with which the cellulose is mixed; and 
when they are removed, the cellulose left behind is the 
same thing, no matter what its source may have been. 
The coarse wood of trees and the tender shoots of the most 
delicate plants consist essentially of cellulose. Cotton- 
wool, hemp, and flax consist almost wholly of cellulose. 

Properties. — Cellulose does not crystallize, and is insol- 
uble in all ordinary solvents. It dissolves in concentrated 
sulphuric acid. If the solution is diluted and boiled, the 
cellulose is converted into dextrin and dextrose. It will 
thus be seen that rags, paper, and wood, all of which con- 



406 INTRODUCTION TO CHEMISTRY. 

sist largely of cellulose, might be used for the preparation 
of dextrose or glucose, and consequently of alcohol. 

Gun-cotton, Pyroxylin, Nitrocellulose. — Cellulose has 
some of the properties of alcohols; among them the power 
to form ethereal salts with acids. Thus, when treated with 
nitric acid it forms several nitrates, just as glycerin forms 
the nitrate known as nitroglycerin (which see). The 
nitrates are explosive, and are used for blasting under the 
name gun-cotton. 

Collodion. — A solution of gun-cotton in a mixture of 
ether and alcohol i3 known as collodion solution, which is 
much used in photography. When poured upon any sur- 
face, such as glass, the ether and alcohol rapidly evaporate, 
leaving a thin coating of gun-cotton. 

Celluloid. — Celluloid is an intimate mixture of gun-cot- 
ton and camphor. As it is plastic at a slightly elevated 
temperature, it can easily be moulded into any desired 
shape. When cooled it hardens. 

Paper. — Paper in its many forms consists mainly of cel- 
lulose. The essential features in the manufacture of paper 
are, first, the disintegration of the substances used. This 
is effected partly mechanically and partly by boiling with 
caustic soda. Then the resulting mass is converted into 
pulp by means of knives placed on rollers. The pulp, 
with the necessary quantity of water, is then passed be- 
tween rollers. Eags of cotton or linen are chiefly used in 
the manufacture of paper; wood and straw 7 are also used. 

Starch, C 6 H 10 O 5 . — Starch is found everywhere in the 
vegetable kingdom in large quantity, particularly in all 
kinds of grain, as maize, wheat, etc.; in tubers, as the po- 
tato, arrowroot, etc. ; in fruits, as chestnuts, acorns, etc. 



STARCH-FLOUR 407 

Manufacture of Starch. — In the United States starch is 
manufactured mainly from maize; in Europe, from pota- 
toes. The processes made use of are mostly mechanical. 
The maize is first treated with warm water; the softened 
grain is then ground between stones, a stream of water 
running constantly into the mill. The thin paste which 
is carried away is brought upon sieves of silk bolting-cloth, 
which are kept in constant motion. The starch passes 
through with the water as a milky fluid. This is allowed 
to settle when the water is drawn off. The starch is next 
treated with water containing a little alkali, the object of 
which is to dissolve gluten, oil, etc. The mixture is now 
brought into shallow, long wooden runs, where the starch 
is deposited, the alkaline water running off. Finally, the 
starch is washed with water, and dried at a low tempera- 
ture. 

Properties. — Starch in its usual condition is insoluble in 
water. If ground with cold water it is partly dissolved. 
If heated with water the membranes of the cells of which 
the starch is composed are broken, and the contents form 
a partial solution. On cooling, it forms a transparent jelly 
called starch-paste. By dilute acids and ferments starch 
is converted into dextrin, maltose, and dextrose. 

Flour. — Wheat flour, which may serve as an example of 
flour in general, contains water, starch with a little sugar 
and gum, gluten, and a small quantity of mineral matter. 
The finest flour contains about 10 per cent of gluten and 
TO per cent of starch. Gluten is a substance that re- 
sembles in many respects the white of eggs, or egg-albu- 
min. 

Bread-making. — The chemical changes which take place 
in bread-making are of special interest. Bread is made by 
mixing the flour with water and a little yeast. The dough 
thus prepared is put in a warm place for a time, when it 



408 INTRODUCTION TO CHEMISTRY. 

rises. The rising is a result of fermentation caused by the 
yeast. A part of the starch contained in the flour is con- 
verted into sugar, and this is then converted into alcohol 
and carbon dioxide by fermentation. The alcohol passes 
off for the most part, and the carbon dioxide in striving to 
escape from the thick gummy dough fills the mass with 
bubbles of gas, making it light and porous. When the 
loaf is put into the oven the gases contained in it ex- 
pand, making it still lighter; then the fermentation is 
checked by the heat and no further chemical change takes 
place except on the surface, where the substances are partly 
decomposed and converted into a dark-colored product, the 
crust. 

A FEW COMPOUNDS FROM COAL-TAR. 

Aromatic Compounds. — The fact that benzene, C 6 H 6 , 
toluene, C 7 H 8 , and other hydrocarbons are obtained from 
coal-tar has already been mentioned (p. 383). These hy- 
drocarbons are the starting-points for the preparation of a 
very large number of compounds of carbon which are com- 
monly called the " aromatic compounds," as many of them 
have a pleasant aromatic odor. 

Nitrobenzene, C f H 5 N0 2 . — This substance is formed by 
treating benzene with nitric acid : 

C 6 H 6 + HN0 3 = C 6 H 5 N0 2 + H 2 0. 

It is a yellow liquid with a pleasant odor like that of the 
oil of bitter almonds. It is much used under the name 
artificial oil of bitter almonds. 

Aniline, C 6 H h NH 2 . — When nitrobenzene is treated with 
a solution from which hydrogen is given off the oxygen is 
extracted and replaced by hydrogen : 

C e H B NO a + 6H = C 6 H 5 NH 2 + 2H 2 0. 



COMPOUNDS FROM COAL TAR. 409 

The product is the substance known as aniline. It is a 
colorless liquid. When it together with a similar substance, 
known as toluidine, is treated with mercuric chloride, 
HgCl 2 , or arsenic acid it is converted into the dye magenta, 
which is the substance from which most of the aniline dyes 
are prepared. 

Aniline Dyes. — Of these a large number are in use. 
They are all derivatives of rosaniline, of which magenta is 
a salt. A great many different colors of aniline dyes are 
made, some of them of great beauty. 

Phenol, Carbolic Acid, C^O. — This familiar substance 
is contained in coal-tar, and is extracted from it by treating 
with caustic soda in which the carbolic acid dissolves. 
When pure it crystallizes in beautiful colorless rhombic 
needles. It has a peculiar, penetrating odor, and is poi- 
sonous. It is much used as a disinfectant. 

Oil of Bitter Almonds, I « „ ~ ml . , . 

_> . A1J , , > C,H fi O. — ihis substance occurs 

Benzoic Aldehyde, ) 7 6 

in combination with amygdalin, which is found in bitter 
almonds, laurel-leaves, cherry-kernels, etc. Amygdalin 
belongs to the class of compounds known as glucosicles, 
which break up into glucose and other substances. Amyg- 
dalin itself, under the influence of emulsin, which occurs 
with it in the plants, breaks up into oil of bitter almonds, 
hydrocyanic acid, and dextrose : 

C 20 H 21 NO U + 2H,0 = C,H.O + CNH + 2C 6 H ]2 6 . 

Amygdalin. Oil of Hydrocy- Glucose, 

bitter anic acid, 

almonds. 

It is prepared from bitter almonds, which yield about 1.5 
to 2 per cent. It is a liquid which has a pleasant odor. 
It is made artificially from coal-tar, and is used in the 
preparation of artificial indigo. 



410 INTRODUCTION TO CHEMISTRY. 

Benzoic Acid, C,H 6 2 . — Benzoic acid occurs in gum ben- 
zoin and in the balsams of Peru and Tolu, and is made 
artificially from coal-tar by oxidizing toluene,* C,H 8 . 

Balsams and Odoriferous Resins. — The balsams of Peru 
and Tolu are thick fragrant fluids which are obtained from 
certain trees in South America and elsewhere by cutting 
the bark. Benzoin is a similar substance. These as well 
as myrrh, frankincense, and other substances of the kind 
are used for their odors. The odors are intensified when 
the substances are heated. They are largely used as incense. 

Gallic Acid, C 7 H 6 6 . — Gallic acid occurs in sumach, in 
Chinese tea, and many other plants. It is formed by boil- 
ing tannin or tannic acid with sulphuric acid. It is pre- 
pared from gall-nuts by fermentation of the tannin con- 
tained in them. It is closely related to tannin or tannic 
acid. 

Tannic Acid, Tannin, C 14 H 10 O 3 .— This substance occurs 
in gall-nuts, from which it is extracted in large quantities. 
It is soluble in water. Its solution gives a dark blue-black 
color with iron salts. Tannin is used extensively in medi- 
cine, in dyeing, in the manufacture of laather and of ink. 

Experiment 193.— Boil 10 grams of powdered gall-nuts with 
60 cc. of water, adding water from time to time. A solution of 
tannin is thus obtained. Filter after standing. In a test-tube 
add to some of this solution a few drops of a solution of copperas 
(ferrous sulphate). A colored precipitate is formed which grad- 
ually changes to black. 

Tanning. — The process of tanning consists in treating 
hides, from which the hair has been removed, with an in- 
fusion of hemlock or oak bark, or of sumach-leaves, in 
which there is tannic acid. The acid combines with certain 

* The name toluene comes from the fact that this hydrocarbon was 
first obtained from the balsam of Tolu. 



COMPOUNDS FROM COAL-TAR— GLUCOSIDES. 41i 

parts of the hides, forming insoluble compounds which 
remain in the pores, converting the hides into leather. 

Indigo. — In several plants which grow in the East and 
West Indies, in South America, Egypt, and other warm 
countries, there occurs a substance called indican which, 
when treated with dilute mineral acids or certain fer- 
ments, breaks up into indigo-blue and a substance resem- 
bling glucose. Commercial indigo contains as its princi- 
pal ingredient indigo-blue. Indigo-blue has been prepared 
artificially from the oil of bitter almonds by the aid of 
complicated processes. 

Naphthalene, C ]0 H 8 . — This hydrocarbon is contained in 
coal-tar in large quantity. It is a beautiful white crystal- 
lized substance much used in the preparation of dyes and 
for protecting 'woollen fabrics from moth. 

Anthracene, C 14 H 10 . — Anthracene like naphthalene is ob- 
tained from coal-tar. Its chief use is in the preparation 
of artificial alizarin. 

Alizarin, C M H fi 4 . — Alizarin is the well-known dye ob- 
tained from madder-root. For some years it has been 
made artificially from anthracene, and the cultivation of 
madder has been given up. Madder-root was used for 
dyeing "Turkey-red." Artificial alizarin is almost exclu- 
sively used for this purpose at present. 



Glucosides. — Glucosides are substances that occur in 
nature in the vegetable kingdom, and that break up under 
the influence of ferments and dilute acids into sugar and 
other compounds. Amygdalin has already been mentioned. 
This breaks up into oil of bitter almonds and dextrose. 
Indican, which yields indigo and dextrose, is another ex- 
ample. 



412 INTRODUCTION TO CHEMISTRY, 

Myronic acid, another glucoside, is found in the form of 
the potassium salt in black mustard-seed. When treated 
with myrosin, which is contained in the aqueous extract of 
white mustard-seed, potassium myronate is converted into 
dextrose and oil of mustard. 

Alkaloids.— These compounds occur in plants, and are 
frequently those parts of the plants which are most active 
when taken into the animal body. They are hence some- 
times called the active principles of plants. Many of these 
substances are used in medicine. They all contain nitro- 
gen and in some respects resemble ammonia. Only a few 
of the more important alkaloids need be mentioned here. 

Quinine. — This valuable alkaloid is obtained from the 
outer bark of certain trees which grow in Peru. The bark 
is known as Peruvian bark. 

Cocaine is found in cocoa-leaves. Its hydrochloric-acid 
salt has recently come into prominence in medicine, owing 
to the fact that a small quantity of its solution placed upon 
the eye or the gums or injected beneath the skin causes 
insensibility to pain. 

Nicotine occurs in tobacco-leaves in combination with 
malic acid. 

Morphine and narcotine are the principal alkaloids found 
in opium, which is the evaporated sap that flows from in- 
cisions in the capsules of the white poppy before they are 
ripe. 



t 



CHAPTER XXXI. 

QUALITATIVE ANALYSIS. 

General. — In order to analyze substances chemists make 
use of reactions such as have been studied in the earlier 
parts of this book. To learn to analyze complicated sub- 
stances, long practice and careful study of a great many 
facts are necessary. But simple substances can be ana- 
lyzed by the aid of such facts as have already been studied. 
It has been seen, for example, that certain chlorides are 
insoluble in water; that certain sulphides are insoluble in 
dilute hydrochloric acid; and that other sulphides which 
are soluble in dilute hydrochloric acid are insoluble in 
neutral or alkaline solutions. Advantage is taken of these 
and other similar facts to classify substances according 
to their reactions. A convenient classification for pur- 
poses of analysis is the following: 

Group I. Metals whose chlorides are insoluble or diffi- 
cultly soluble in water. This group includes: Silver, 
lead, and mercury (as mercurous salt). 

Group II. Metals not included in Group I, whose sul- 
phides are, however, insoluble in dilute hydrochloric or 
nitric acid. This group includes: Copper, mercury (as 
mercuric salt), bismuth, antimony, arsenic, and tin. 

Group III. Metals not included in Groups I and II, whose 
sulphides are, however, precipitated by ammonium sul- 
phide and ammonia. This group includes: Alu- 
minium, chromium, nickel, cobalt, iron, zinc, and 
manganese. 

413 



414 INTRODUCTION TO CHEMISTRY. 

Group IV. Metals not included in Groups I, II, and III, 
but which are precipitated by ammonium carbonate, 
ammonia, and ammonium chloride. This group in- 
cludes: Barium, strontium, and calcium. 

Group V. Metals not included in Groups I, II, III, and 
IV, but which are precipitated by disodium phosphate, 
HNa 2 P0 4 , ammonia, and ammonium chloride. This 
group includes : Magnesium. 

Group VI. Metals not included in Groups I, II, III, IV, 
and V. This group includes: Sodium, potassium, 
and ammonium. 

1. Now, suppose you have a substance given you for 
analysis. The first thing to do is to get the substance in 
solution. See whether it dissolves in water. If it does 
not, try dilute hydrochloric acid. If it does not dissolve in 
hydrochloric acid, try nitric acid; and if it does not dis- 
solve in nitric acid, try a mixture of nitric and hydro- 
chloric acids. If concentrated acid is used, evaporate to 
dryness on a water-bath «bef ore proceeding further. Then 
dissolve in water, and add a few drops of hydrochloric 
acid. If a precipitate is formed, continue to add the acid 
drop by drop until a precipitate is no longer formed. 
Filter and wash. 

What may this precipitate contain ? 

2. Pass hydrogen sulphide through the filtrate for some 
time and let stand. Filter and wash. 

If a precipitate is formed, what may it contain ? 

3. Add ammonia and ammonium sulphide to the fil- 
trate. Filter and wash. 

If a precipitate is formed, what may it contain ? 

4. Add ammonium carbonate, ammonia, and ammonium 
chloride to the filtrate. Filter and wash. 

If a precipitate is formed, what may it contain ? 

5. Add disodium phosphate, ammonia, and ammonium 
chloride to the filtrate. Filter and wash. 



EXAMPLES FOR PRACTICE. 415 

If a precipitate is formed, what may it contain ? 
What may be in the filtrate ? 

Examples for Practice. — Before attempting anything in 
the way of systematic analysis it will be well to experiment 
in a more general way, with the object of determining which 
one of a given list of substances a certain specimen is. 

The list below contains the names of the principal sub- 
stances with which you have thus far had directly to deal 
in your work. You have handled them and have seen how 
they act toward different substances. Suppose now that a 
substance is given you, and you know simply that it is one 
of those named in the list, how would you go to work to 
find out which one it is ? You have a right to judge by 
anything in the appearance or in the conduct of the sub- 
stance. If you reach a conclusion, see whether you are 
right by further experiments. After your work is finished 
write out a clear account of what you have done, and state 
your reasons for the conclusion you have reached. 

For example, suppose sodium chloride is given you. 
You see that it is a white solid. On heating it in a small 
tube, you see that it does not melt, but it breaks up into 
smaller pieces with a crackling sound. It is soluble in 
water. Hydrochloric acid causes no change when added 
to a little of the solid. Is it a carbonate ? Sulphuric 
acid causes evolution of a gas. Has this an odor ? How 
does it appear when allowed to escape into the air ? Is it 
nitric acid? Collect some of it in water. How does this 
solution act on a solution of silver nitrate ? By this time 
you have evidence that you are dealing with a chloride, 
but you do not yet know which chloride it is. It cannot 
be ammonium chloride. Why ? It may be either potassium 
or sodium chloride. Try a small piece in the flame. What 
color ? You now have good reasons for believing that the 
substance you are dealing with is sodium chloride. To 
convince yourself, get a small piece of sodium chloride 



416 



INTRODUCTION TO CHEMISTRY. 



from the bottle known to contain it, and make a series of 
parallel experiments with this and see whether you get ex- 
actly the same results that you got with the specimen you 
were examining. If not, account for the differences. 

By careful work there will be no serious difficulty in 
determining which one of the substances in the list you 
are dealing with. 

List of Substances for Examination. 



1. 


Sugar. 


19. 


Manganese dioxide. 


2. 


Mercuric oxide. 


20. 


Charcoal. 


3. 


Calc spar. 


21. 


Calcium sulphate (Gyp 


4. 


Marble. 




sum). 


5. 


Copper. 


22. 


Copper oxide. 


6. 


Hydrochloric acid. 


23. 


Ammonium chloride. 


7. 


Nitric acid. 


24. 


Calcium oxide (Quick 


8. 


Sulphuric acid. 




lime). 


9. 


Zinc. 


25. 


Sodium nitrate. 


10. 


Tin. 


26. 


Ammonium nitrate. 


11. 


Oxalic acid. 


27. 


Sodium chloride. 


12. 


Sodium carbonate. 


28. 


Potassium bromide. 


13. 


Ferrous sulphate (Cop- 29. 


Potassium iodide. 




peras). 


30. 


Iron sulphide. 


14. 


Roll-sulphur. 


31. 


Potassium carbonate. 


15. 


Iron-filings. 


32. 


Potassium nitrate. 


16. 


Carbon bisulphide. 


33. 


Potassium dichromate. 


17. 


Lead. 


34. 


Red lead (Minium). 


18. 


Potassium chlorate. 


35. 


Lead nitrate. 



[The instructor will, of course, select the substance and 
give it to the student without any suggestion as to what it 
is. After the student has shown that he can tell with cer- 
tainty which substance he has, some simple mixtures of 
substances selected from the above list may next be given 
for examination. Thus charcoal and copper oxide; zinc 
and tin; mercuric oxide and iron-filings; etc., etc.] 



GROUPS I AND II. 417 



Study of Gkoup I. 

Experiment 194.— 1. Prepare dilute solutions of silver nitrate, 
AgN0 3 , lead nitrate, Pb(N0 3 ) 2 , and mercurous nitrate, HgNOs. 

2. Add to a small quantity of each separately in test-tubes a 
little hydrochloric acid. 

What is formed ? 

3. Heat each tube with contents, and then let cool. 
What difference do you observe ? 

4. After cooling, add a little ammonia to the contents of each 
tube. 

What takes place in each case ? 

How could you distinguish between silver, lead, and mercury? 

5. Mix the solutions of silver nitrate, lead nitrate, and mer- 
curous nitrate, and add a little of the mixture to considerable 
water in a test-tube. Add hydrochloric acid as long as it causes 
the formation of a precipitate. Heat to boiling. Filter rapidly 
and wash with hot water. 

What is in the filtrate, and what is on the filter ? 

6. Let the filtrate cool. 

What evidence have you that there is anything present in it? 

7. Add sulphuric acid to a little of the liquid. 

8. Add hydrogen sulphide to a little of the liquid. 

9. Pour ammonia on the filter, and wash out with water. Then 
add nitric acid to the filtrate. 

What evidence do you get of the presence of silver and of 
mercury ? 

Study of Group II. 

Experiment 195. — 1. Prepare dilute solutions of copper sul- 
phate, mercuric chloride, arsenic trioxide in hydrochloric acid, 
and of tin in hydrochloric acid. [Bismuth and antimony are 
omitted, as their presence gives rise to difficulties hard to deal 
with intelligently at this stage.] Add a little hydrochloric acid 
to the solutions of copper sulphate and of mercuric chloride. 

2. Pass hydrogen sulphide through a small quantity of each of 
the solutions. 

What takes place ? What are the substances formed ? 



418 INTRODUCTION TO CHEMISTRY. 

3. Filter and wash. Treat each precipitate with a solution of 
yellow ammonium sulphide. 

What takes place ? Add dilute sulphuric acid to the filtrates. 
What takes place ? 

4. Treat the precipitates obtained from the copper and the 
mercury salts with concentrated warm nitric acid. 

Does either one dissolve easily ? What is the color of the 
solution ? 

5. Treat a little of the solution obtained in 4. with ammonia. 
What is the result ? How can you detect the presence of 

copper ? 

6. Treat with a mixture of nitric and hydrochloric acids the 
precipitate which is not readily dissolved by nitric acid alone. 
Evaporate the acid. Add water, and then a solution of tin in 
hydrochloric acid. 

What is formed when tin is dissolved in hydrochloric acid ? 

What other compound of tin and chlorine is there ? 

[When stannous chloride, SnCl 2 , acts upon mercuric chloride, 
HgCl 2 , the former takes a part or all of the chlorine from the 
latter, forming either mercurous chloride, HgCl, or mercury, 
thus : 

2HgCl 2 + SnCl 2 = 2HgCl + SnCU; 
HgCl 2 + SnCl 2 = Hg + SnCU.] 

7. Treat the precipitate obtained in the case of the arsenic with 
4-5 cc. of a concentrated solution of ammonium carbonate. 
To the solution add hydrochloric acid and a few crystals of 
potassium chlorate, and boil until chlorine is no longer given off. 
Add ammonia, ammonium chloride, and magnesium sulphate to 
the solution. [The precipitate is ammonium magnesium arsenate, 
NH 4 MgAs0 4 .] 

8. Dissolve the tin precipitate in dilute hydrochloric acid. 
Add a few small pieces of zinc. Dissolve in hydrochloric acid 
the tin which separates. 

What will the solution thus obtained contain ? 
What should take place on adding the solution to a solution of 
mercuric chloride ? Try it. 

Mix the solutions prepared in 1., and proceed as follows : 

9. Pass hydrogen sulphide. Filter; wash. Treat the precipitate 
with ammonium sulphide. Filter; wash. 



GROUP III. 419 

What is now in solution ? 
What is on the filter ? 

10. Treat the solution with dilute sulphuric acid. Filter ; 
wash. Treat the precipitate thus obtained with concentrated 
ammonium carbonate. Filter ; wash. Treat the solution as 
directed in 7., and the precipitate as in 8. 

11. Treat with concentrated warm nitric and hydrochloric 
acids the precipitate left after treating with ammonium sulphide 
as in 9. Test for copper as in 5., and for mercury as in 6. 



Study of Group III. 

ALUMINIUM. 

Experiment 196.— 1. Prepare a solution of ordinary alum. 
[What is ordinary alum ?] 

2. Add to this solution ammonia, ammonium chloride, and 
ammonium sulphide. Filter and wash. Treat the precipitate 
with hydrochloric acid ; and then treat the solution thus obtained 
with ammonium chloride and ammonia. 

[Aluminium does not form a sulphide ; but the hydroxide, 
Al(OH) 3 , is formed when ammonia, ammonium chloride, and 
ammonium sulphide are added to a solution of its salts. When 
the hydroxide is treated with hydrochloric acid it is converted 
into the chloride, AlCl a , which dissolves ; and when the solution 
of the chloride is treated with ammonia the hydroxide is pre- 
cipitated : 

Aids + 3NH 3 + 3H 2 = Al(OH) s 4- 3NH*C1.] 

3. Dissolve the precipitate of aluminium hydroxide, Al(OH) : , 
in as little hydrochloric acid as possible, and add a cold solutio.. 
of sodium hydroxide. Boil the solution thus obtained. 

4. After cooling slowly add dilute hydrochloric acid. When 
the alkali is neutralized, aluminium hydroxide, Al(OH) 3 , will be 
precipitated. It will dissolve on the addition of more acid ; and 
from the solution thus obtained the hydroxide can be precipitated 
by a solution of ammonia. 



420 INTRODUCTION TO CHEMISTRY. 

CHROMIUM. 

Experiment 197.— 1. To 5-10 cc. of a solution of potassium 
dichromate in a test-tube add 10-15 drops of hydrochloric acid 
and 10-15 drops of alcohol, and boil. What change takes 
place ? 

[Under the conditions the chromium is changed to chromic 
chloride, CrCU , and the potassium to potassium chloride, while 
some of the oxygen of the dichromate acts upon the alcohol, con- 
verting it into aldehyde : 

K 2 Cr 2 7 + 8HC1 = 2KC1 + 2CrCl 3 4- 4H 2 + 30 ; 
3C 2 H 6 + 30 = 3C 2 H 4 + 3H 2 0.j 

Alcohol. Aldehyde. 

2. Treat the solution of chromic chloride, CrCl 3 , obtained in 1. 
as directed in 2. and 3., Experiment 196, and note the differ- 
ences. 

How could you distinguish between aluminium and chromium ? 

IRON. 

Experiment 198. — 1. Prepare a solution containing ferrous 
chloride. [See Experiment 180.] 

2. Convert a part of this into ferric chloride. [See Experi- 
ment 180.] 

3. Treat each of these solutions with ammonia and ammonium 
sulphide. 

[The precipitate is the same in both cases, and the action is 
represented thus : 

FeCl 2 + (NH 4 ) 2 S = FeS + 2OTLC1 ; 

2FeCl 3 + 3(NH 4 ) 2 S = 2FeS + 6NH 4 C1 + S.] 

4. Dissolve the precipitate in hydrochloric acid : 

FeS + 2HC1 = FeCl 2 + H 2 S. 

5. Convert the ferrous into ferric chloride. [See Experiment 
180.] 

6. Treat with ammonium chloride and ammonia. Filter and 
wash. Treat the precipitate as directed under 3., Experiment 
196, 



ORG UP III. 421 

What differences are there between aluminium, chromium, and 
iron ? 

7. Filter ; dissolve the precipitate in hydrochloric acid ; and 
treat with a solution of potassium ferrocyanide, K 4 Fe(CN) 6 . 

The precipitate formed in this case is Prussian blue. 

ZINC. 

Experiment 199.— 1. Prepare a dilute solution of zinc sul- 
phate. 

2. Treat with ammonia and ammonium sulphide. What is the 
color of the precipitate ? The composition is ZnS. 

3. Dissolve in dilute hydrochloric acid : 

ZnS + 2HC1 = ZnCl 2 + H 2 S. 

4. Treat with ammonium chloride and ammonia. Is a pre- 
cipitate formed ? 

5. Add enough hydrochloric acid to give the solution an acid 
reaction, and then add sodium acetate, NaC 2 H 3 2 : 

ZnCl 2 + 2NaC 2 H 3 2 = 2NaCl + Zn(C 2 H 3 2 ) 2 . 

6. Pass hydrogen sulphide through the solution. The white 
precipitate is zinc sulphide, ZnS. 

What differences are there between aluminium, chromium, 
iron, and zinc ? How could they be separated and detected if 
present in the same solution ? 

[It will be well for the instructor to prepare solutions contain- 
ing two or more members of Group III, and to give them to the 
student for analysis.] 

MANGANESE. 

Experiment 200.— 1. Treat a little manganese dioxide in a 
test-tube with hydrochloric acid. Boil, dilute, and filter. 
What have you in solution ? [See page 97.] 

2. Treat as under 2., 3., 4., 5., 6., in the preceding Experiment. 
In what respects do manganese and zinc differ ? 

3. To the solution through which you have just passed hydro- 
gen sulphide add sodium hydroxide, NaOH, until most of the 
acetic acid is neutralized ; heat gently and add bromine- water. 
Let the liquid stand for an hour. 



422 



INTRODUCTION TO CHEMISTRY. 



What takes place ? [The composition of the precipitate is 
represented by the formula Mn(OH) 4 .] 

How could you separate manganese from the other members 
of the group ? 

Experiment 201. — 1. Mix dilute solutions of alum, chromic 
chloride (prepared as in Experiment 197, 1.), ferrous chloride 
(prepared as in Experiment 180), zinc sulphate, and manganous 
chloride. 

2. Treat with ammonia, ammonium chloride, and ammonium 
sulphide. Filter and wash. 

3. Treat the precipitate with dilute hydrochloric acid ; treat 
with nitric acid to convert ferrous chloride into ferric chloride 
(Experiment 180) ; and then treat the solution thus obtained 
with ammonium chloride and ammonia. 

What have you in the precipitate ? (Call this A.) 
What in the solution ? (Call this B.) 

4. Dissolve the precipitate in a little dilute hydrochloric acid, 
and add a cold solution of sodium hydroxide, more than enough 
to neutralize the hydrochloric acicl. Filter ; dissolve the precipi- 
tate in hydrochloric acid ; and treat with a solution of potassium 
ferrocyanide, K 4 Fe(CN) 6 . [See Experiment 198, 7.] Boil the 
nitrate from the precipitate of ferric hydroxide. What is pre- 
cipitated ? Treat the nitrate as directed in Experiment 196, 4. 

5. Treat the solution B (see under 3., above) as directed under 
5. and 6., Experiment 199 ; and under 3., Experiment 200. Ex- 
amine mixtures containing members of Group III. 



Study of Group IV. 



CALCIUM. 



Experiment 202. — 1. Prepare a solution of calcium chloride by 
dissolving a little calcium carbonate (marble) in hydrochloric 
acid. What is the reaction ? 

2. Treat with ammonium chloride, ammonia, and ammonium 
carbonate, (ISTH^COa. Filter and wash. 

What takes place ? Write the equation. 

3. Dissolve the precipitate in dilute hydrochloric acid. Treat 
a small part of this solution with a solution of calcium sulphate 



GROUPS IV AND r. 423 

in water. Treat another small pari with ammonia and ammonium 
oxalate, ( N 1 1 4 >,<.<> 4 . The precipitate It calcium oxalate, OaOgOi 
Docs a solution of calcium chloride give a precipitate when 
( reated wit h a solut ion of calcium sulphate I 



BARIUM. 

Experiment 208, 1. Prepare a dilute solution of barium ohlo- 
ride in water. 

2. Treat as directed under 2., preceding Experiment, 

\\. Dissolve the precipitate in dilute hydrochloric acid. Treat a 
small part of this solution with a solution of calcium sulphate In 
water. 

What difference do you notice between the oonduot of calcium 
and that of barium i 

How could you detect barium and calcium when present In the 
same solution! Mix the solutions of barium and calcium chlo- 
rides, and try the reactions described in Experiments 209 and 
208. 



Study OF GROUP V. 

MAGNESIUM. 

KXPERIMKNT 20-1. 1. Prepare a dilute solution of magnesium 
Sulphate in water. 

2. Add ammonium chloride, ammonia, and disodium phos- 
phate, IINaJN),. 

The precipitate formed is ammonium magnesium phosphate, 
N II.M^PO,. What similar precipitate has already been obtained i 

(See Experiment 195, 7.) 

:*. Mix solutions of barium chloride, calcium chloride, and 

magnesium chloride; and see whether you can detect the three 
metals by means of the reactions described in Experiments 202, 

208, and 204. 



424 INTRODUCTION TO CHEMISTRY. 



Study of Group VI. 

Experiment 205. — 1. Potassium can be detected by means of 
the color it gives to a flame (see Experiment 149) ; and also by the 
fact that when platinic chloride, PtCl 4 , is added to a solution of 
a potassium salt, the salt, K 2 PtCl 6 , is precipitated. Try this. 

2. Sodium is detected by means of the flame reaction (see Ex- 
periment 149.) 

3. Ammonium salts are detected by adding an alkali, when 
ammonia gas is given off, and this is easily recognized. 

General Directions. — By the aid of the reactions thus far 
studied it will be found possible to analyze substances con- 
taining the following metals either alone or mixed together: 
Silver, lead, mercury, copper, tin, arsenic, aluminium, 
chromium, iron, zinc, manganese, calcium, barium, mag- 
nesium, potassium, sodium, and ammonium. After the 
metals have been detected, the next question to be an- 
swered is : In what forms of combination were they present 
in the original substance taken for analysis ? Or, in other 
words, what salts were present ? To answer this question, 
recall the experiments you have made in the general re- 
actions of chlorides, nitrates, sulphates, and carbonates. 
These are the most common salts and, for the present, it 
will be best to confine your work to these. 

Classification of Substances Studied. — It will now be well 
to draw up a table containing the names and symbols of 
all the substances with which you have had to deal, classi- 
fying them into : 

(1) Elements and Compounds ; 

(2) Acids, Bases, and Salts. 

Under Elements state the principal source and the prin- 
cipal method of getting each. 

Under Compounds state the source and the principal 
method of preparation of each. 



GROUP VL 425 

Classify all the compounds you have had to deal with 
into : 

(1) Those which are gaseous ; 

(2) Those which are liquid ; 

(3) Those which are solid at the ordinary temperature; 

(4) Those solids and liquids which easily undergo 
change when heated ; (state what the change is, and give 
the equation expressing the change.) 

Classify the compounds further into : 

(1) Those which are soluble in water without change; 

(2) Those which dissolve in water and are changed; 
(state what the change is, and give the equation express- 
ing the change.) 

(3) Those which are insoluble in water. 



INDEX. 



Acetates, 393 
Acetylene, 386, 388 
Acid, acetic, 383, 393 

arsenic, 251 

arsenious, 251 

benzoic, 410 

boric, 256 

bromic, 220 

butyric, 394 

carbolic, 409 

carbonic, 176 

chloric, 111, 115 

chlorous. 115 

chromic, 372 

citric, 397 

dithionic, 241 

ferric, 367 

formic, 393 

gallic, 410 

hydriodic, 222 

hydrobromic, 219 

hydrochloric, 40, 95, 96, 97, 
103, 104 

hydrocyanic, 190 

hydrofluoric, 223 

hypobromous, 220 

hypochlorous, 115 

hyposulphurous, 241 

iodic, 222 

lactic, 396 

malic, 397 

metaboric, 256 

metaphosphoric, 248 

metarsenic, 251 

metastannic, 356 

myronic, 412 

nitric, 14, 15, 40, 134, 143 

nitrosyl-sulphuric, 238 



Acid, nitrous, 134, 148 

oleic, 398 

orthophosphoric, 247 

oxalic, 178, 396 

palmitic, 394 

perchloric, 115 

phosphoric, 247 

phosphorous, 248 

propionic, 394 

prussic, 190 

pyroarsenic, 251 

pyrogallic, 335 

pyrophosphoric, 248 

pyrosulphuric, 241 

salts, 241 

selenic, 242 

silicic, 259 

stannic, 355 

stearic, 394 

sulphuric, 14, 40, 95, 97, 104, 
130, 237 

sulphurous, 236 

tannic, 410 

tartaric, 397 

telluric, 242 

tetraboric, 256 

tetrathionic, 241 

thiosulphuric, 241 

trithionic, 241 
Acid- forming elements, 208 
Acids, 40, 91, 116, 120, 208 

dibasic, 240 

fatty, 394 

monobasic, 240 

organic, 401 

tribasic, 247 
Agate, 260 
Air, 125, 217 

427 



428 



INDEX. 



Alabaster, 310 
Albumin, 157, 382 
Alcohol, ethyl, 383, 389 

methyl, 383, 389 
Alcohols, 400 
Aldehyde, 392 

benzoic, 409 
Alizarin, 411 
Alkalies, 116, 283 
Alkaloids, 412 
Allylene, 386 
Alum, ordinary, 343 
Aluminates, 342 
Aluminium, 339 

alloys, 340, 342 

chloride, 340 

compounds, 342, 347, 419 

group, 339 

hydroxide, 342 

oxide, 342 

silicates, 343 

sulphate, 342 
Alums, 343 

Amalgamation -process, 333 
Amalgams, 330 
Amethyst, 260 
Ammonia, 133, 134, 141, 148 

composition, 139 

formation. 133 

in the air, 132 

water, 162 
Ammonium, 137 

chloride, 134, 294 

chlorplatinate, 377 

compounds, 424 

hydrosulphide, 298 

hydroxide, 138 

nitrate, 151 

salts, 138, 296 

sodium phosphate, 299 

sulphide, 297 
Amygdalin, 409 
Anaesthetics, 388, 397 
Analysis, 49, 234, 274, 413 

of water, 50 
Anhydride, boric, 257 

carbonic, 170 

nitric, 150 

nitrous, 150 

permanganic, 369 

silicic, 260 



Anhydrides, 150 
Anhydrite, 310 
Aniline, 408 

dyes, 409 
Anthracene, 383, 411 
Anthracite Coal, 162 
Antimony, 98, 251 

acids of, 252 

salts of, 252 

trichloride, 99 
Apatite, 243, 247, 312 
Aqua regia, 148 
Aragonite, 308 
Aromatic compounds, 408 
Arsenic, 248 

acids of, 251 

compounds, 417 

trioxide, 248, 250 

white, 166 
Arsine, 249 
Asbestos, 319 
Ash, black, 294 
Atomic theory, 77 

weights, 79, 214 
determination of, 80, 196, 
336 
Atoms, 77 
Avogadro's Law, 194 

Balloons, 44 
Balsams, 410 
Barium, 317 

compounds, 423 

dioxide, 67, 93, 318 

hydroxide, 318 

oxide, 317 

sulphate, 277 
Baryta- water, 129, 130, 165, 318 
Base-forming elements, 208 
Bases, 116, 120, 122, 208, 261 
Bauxite, 339 
Bell-metal, 327 
Benzene, 383, 386, 408 
Benzine, 385 
Benzoin, 410 

Beryllium (= glucinum), 303 
Bessemer steel, 362 
Bismuth, 253 

salts of, 253 
Bituminous coal, 162 
Blast-furnace, 359 



1KDEX. 



429 



Bleaching, 99, 100, 236 

powder, 112, 113, 306 
Blowpipe, 185 
Bone-black, 160 
Bone-oil, 383 
Boracite, 256 
Borax, 256, 296 
Boric acid, 256 

anhydride, 257 
Boron, 255 

chloride, 256 

crystallized, 255 
Brass, 323, 326 
Bread-making, 407 
Breathing, 34 
Bricks, 347 
Britannia metal, 354 
Bromides, 218 
Bromine, 217 
Bronze, 327 
Bunsen burner, 188 
Burning in the air, 31 
Butane, 385 
Butter, 399 
Butylene, 386 

Cadmium, 319 
Caesium, 300 
Calcite, 308 
Calcium, 303 

carbonate, 308 

chloride, 49, 109, 129, 130, 135, 
304 

compounds, 303, 422 

hydroxide, 305 

hypochlorite, 306 

s:roup, 303 

oxide, 304 

phosphates, 312 

silicates, 314 

sulphate, 310 

sulphide, 317 
Calomel, 331 

Cane-sugar, 402, 403, 404 
Caramel, 404 
Carbohydrates, 402 
Carbon, 26, 88, 91, 156, 165 

bisulphide, 230, 241, 387 

dioxide, 168 
cycle of in nature, 174 
in the air, 129, 132 



Carbon group, 258, 260 

monoxide, 177 
Carbonates, 169, 176, 279 
Carnelian, 260 
Casein, 405 
Cassiterite, 353 
Cast-iron, 360 
Caustic potash, 38, 110, 2*5 

soda, 38, 130, 291 
Celluloid, 406 
Cellulose, 158, 382, 405 
Cementation, 361 
Cemeuts, 317 
Cerium, 260 
Cerussite, 353 
Chalcocite, 325 
Chalk, 308 
Charcoal, 158 

animal, 160 

filters, 160 

reduction by, 166 

wood, 158, 159 
Chemical action, 12, 15 

changes, 2 

energy, 36 

work, 36 
Chemistry, 2 
Chloral, 392 

hydrate, 392 
Chlorates, 276 
Chlorides, 102, 266, 268 
Chlorine, 95, 100, 217 

acids, 115 

bleaching by, 100 

comparison with fluorine, bro- 
mine, and iodine 217 

hydrate, 101 

oxides of, 115 
Chloroform, 388 
Chromates, 372, 373 
Chrome alum, 374 

yellow, 352, 374 
Chromic acid, 372 

chloride, 373, 420 

iron, 372 
Chromium, 372 

compounds, 372, 420 
Cinnabar, 330 
Clay, 339, 344 
Coal, 162, 383 
Coal-gas, 180 



430 



INDEX. 



Coal-tar, 162, 383, 408 

Cobalt, 367 

Cocaine, 412 

Coke, 160 

Collection of gases, 8 

Collodion, 406 

Columbium, 254 

CombiDation, laws of chemical, 

71, 73 
Combining-weights, 28, 30 
CombustioD, 32, 169 
Compounds, chemical, 12 
Conservation of energy, 71, 175 
Copper, 14, 99, 152, 325, 387 

acetate, 393 

alloys, 326 

compounds, 327, 417 

group, 325 

oxide, 93, 166, 329 

pyrites, 325 

sulphate, 96, 329 

sulphide, 329 
Copper-plating, 329 
Copperas, 366 
Corrosive sublimate, 331 , 
Corundum, 342 
Cryolite, 223, 289, 339 
Crystallography, 228 
Cupric compounds, 328 
Cuprous compounds, 328 

oxide, 328 
Cyanogen, 189 

Deacon's Process, 96 
Decomposition, double, 93 

heat of, 35 
Decrepitation, 291 
Definite proportions, law of, 71 
Deliquescence, 49 
"Developers," 335 
Dextrin, 402 
Dextrose, 402, 403, 409 
Diamond, 157 
Distillation, 64 
Dolomite, 168, 319 
Double decomposition, 93 
Dulong and Petit's Law, 338 
Dynamite, 399 

Earthenware, 346 
Efflorescence, 48 



Electric currents, 6 
Electrotype, 330 
Elements, 10, 18, 19, 197, 204, 
209, 212 

base-forming, 261 

list of, 19, 212 

molecules of, 197 

names of, 19 

natural families of, 209 

specific heat of, 336 

substituting power of, 204 
Emery, 342 
Emulsin, 409 
Epsom salt, 321 
Erbium, 19 
Essences, 399 
Etching, 224 
Ethane, 385 
Ether, 397 
Etheral salts, 397 
Ethers, 397 
Ethyl alcohol, 389 

butyrate, 399 

nitrate, 398 
Ethylene, 386, 388 
Eudiometer, 53 

Explosion of hydrogen and oxy- 
gen, 53, 59 

Fats, 398 

Feldspar, 283, 339, 343, 344 
Fermentation, 169, 383, 389, 390 
Ferric acid, 367 

chloride, 365 

compounds, 365 

hydroxide, 366 

oxide, 366 

sulphate, 364 
Ferrous chloride, 365 

compounds, 365 

h}'droxide, 365 

oxide, 366 

sulphate, 335, 366 
Fertilizers, 313 
Fibrin, 382 
Fire-damp, 386 
Flame, 180 

reactions, 300 
Flames, 180, 181, 184 

luminosity of, 187 
Flour, 407 



INDEX. 



431 



Fluorine, 223 
Fluor-spar, 223 
Formulas, 81, 201 
Formulas, chemical, 81 
Frankliuite, 321 
Fusel-oil, 391 

Gahnite, 321 
Galenite, 349 
Gall-nuts, 410 
Gallium, 214, 348 
Galvanized iron, 323 
Gas, illuminating-. 180 

marsh-, 385, 386 

olefiant, 388 

water-, 40, 91, 179 
Gases, combination by volume, 
140 

measurement of, 54-59 

specific gravity of, 142, 194 

volumes of combining, 140, 200 
Gasoline, 385 
Gasometer, 23 
German silver, 327, 367 
Germanium, 214, 349 
Glass, 260, 314 
Glucinum, 303 
Glucose, 389, 402 
Glucosides, 409, 411 
Glycerin, 391, 400 
Gold, 376, 378, 380 

alloys, 380 

chlorides, 380, 381 

mining, 379 
Granite, 344 
Grape-sugar, 389 
Graphite, 158 
Guano, 396 
Gun-cotton, 406 
Gun-metal, 327 
Gunpowder, 287 
Gypsum, 310 

Haematite, 358, 366 

Hard water, 177, 309, 311, 395 

Heat of combustion, 34 

decomposition, 35 
Heptane, 385 
Hexane, 385 
Homologous series, 385 
Homology, 385 



Hornblende, 319 
Hydrocarbons, 168, 384 
Hydrogen, 17, 37, 89 

dioxide, 67. 93 

sulphide, 231 , 387 
Hydrosulphides, 234 
Hydroxides, 269 

Iceland spar, 308 
Illuminating-gas, 180 
Illumination, 180 
Incense, 410 

Indestructibility of matter, 70 
Indican, 411 
Indigo, 411 
Indium, 339 
Ink, 410 

sympathetic, 368 
Invert-sugar, 404 
Iodine, 220 
Iodoform, 388 
Iridium, 376, 377 
Iron, 39, 90, 326, 358, 363 

alum, 366 

compounds, 364, 420 

galvanized, 323 

group, 358 

passive, 364 

pyrites, 367 

rust, 363 

sulphides, 367 

varieties of, 360 

Kainite, 289 

Kaolin, 339, 344 

Kelp, 221 

Kerosene, 385 

Kieserite, 321 

Kindling-temperature, 32, 182 

Lamp-black, 160 

Lanthanum, 339 

Lapis lazuli, 345 

"Laughing-gas," 151 

Law of Avagadro, 194 
definite proportions, 71 
Dulong and Petit, 338 
multiple proportions, 73, 74 
specific gravities of gases, 194 
heats of elements, 338 

Lead, 349 



432 



INDEX. 



Lead acetate, 352, 393 

alloys, 351, 354 

black-, 158 

carbonate, 353 

chloride, 352 

chromate, 352, 374 

oxide, 351 

peroxide, 352 

salts, 352, 417 

sulphate, 352 

sulphide, 353 

white-, 353 
Leather, preparation of, 410 
Le Blanc process, 293 
Lepidolite, 300 
Levulose, 402, 403 
Light, chemical action of, 103, 

335 
Lignite, 162 
Lime, 109, 135, 304 

light, 63 
Lime-water, 129, 130, 305 
Litharge, 351 
Lithium, 300 
Litmus, 107, 116 
Luminosity of flames, 187 
" Lunar caustic/' 334 

Madder-root, 411 
Magenta, 409 
Magnesia, 109, 320 
Magnesite, 319 
Magnesium, 15, 319 

carbonate, 320 

chloride, 320 

compounds, 320, 423 

oxide, 109, 320 

sulphate, 321 
Magnetite, 358 
Manganese chloride, 97 

compounds, 369, 421 

dioxide, 22, 87, 96, 369 

group, 369 

salts, 369 

sulphate, 97, 219 
Marble, 13, 303, 308 
Marl, 309, 345 
Marsh-gas, 385, 386 
Marsh's test, 250 
Martin steel, 362 
Matches, 245, 288 



Measurement of gases, 54-59 

heat, 35 
Mechanical mixtures, 10 
Meudeleeff's periodic law, 210 

scheme, 212 
Meerschaum, 319 
Mercuric chloride, 331 

oxide, 20, 83, 331 

sulphide, 331 
Mercurous chloride, 331 
Mercury, 330 

compounds, 330, 417 
Metallic properties, 121 
Metals, 263 
Metathesis, 93 
Methane, 385 
Mica, 339, 344 
Microcosmic salt, 299 
Milk-sugar, 404 
Minium, 351 

Mixtures, mechanical, 10 
Molasses, 404 
Molecular formulas, 201 

weights, determination of, 195 
Molecules, 81, 195, 200 

of the elements, 197 
Molybdenum, 19 
Morphine, 412 
Mortar, 316 

Multiple proportions, law of, 73, 
74 

Naphtha, 385 
Naphthalene, 383, 411 
Narcotine, 412 
Nascent state, 199 
Natural waters, 63 
Neutralization, 116 
Nickel, 367 
Nicotine, 412 
Nitrates, 148, 275 
Nitric oxide, 147, 152 
Nitrification, 134, 143 
Nitrobenzene, 408 
Nitrocellulose, 406 
Nitroglycerin, 399 
Nitrogen, 126, 243 

group, 243, 254 

oxides of, 150 

pentoxide, 150 

peroxide, 147, 153 



INDEX. 



433 



Nitrogen trioxide, 149 
Nitrous anhydride, 149 

oxide, 151 
Nomenclature of acids, 121 

bases, 122 

chlorides, 102 

oxides, 102 

salts, 122 

Octane, 385 

Oil of bitter almonds, 409 

illuminating, 384 

of vitriol, 239 
Oleomargarin, 399 
Opium, 412 
Ores, 264 
Osmium, 376 
Oxidation, 62, 187 

slow, 33 
Oxides, 36, 269 

of nitrogen, 150 
uses of, 154 
Oxygen, 17, 20, 24 

and the sulphur group, 242 
Oxyhydrogen blowpipe, 62 
Ozone, 67, 198 

Palladium, 376 

Paper, 406 

Paraffin, 385 

Paraldehyde, 392 

Passive iron, 364 

Pattinson's method, 332 

Peat, 162 

Pentane, 385 

Periodic law, 210 

Peruvian bark, 412 

Petroleum, 384 

Phenol, 409 

Phosphates, 280 

Phosphine, 245 

Phosphonium salts, 246 

Phosphorescence; 317 

Phosphorite, 243, 247, 312 

Phosphorus, 25, 27, 88, 125, 243 

acids of, 247 

pentoxide, 89, 247 

red, 244 
Photography, 335 
Photometer, 181 
Physical changes, 2 



Pinchbeck, 327 
Pitchblende, 375 
Plaster of Paris, 310 
Platinic chloride, 378 
Platinous chloride, 378 
Platinum, 376 

alloys, 377 

bases, 378 
Plumbago, 158 
Porcelain, 344, 345 
Potash, caustic, 38, 110, 285 
Potassium, 37, 90, 108, 283 

acid tartrate, 397 

chlorate, 21, 84, 110, 112, 288 

chloride, 87, 110 

chlorplatinate, 299, 378 

chromate, 372 

compounds, 284, 424 

cyanide, 189, 190, 288 

dichromate, 372 

ferrocyanide, 190 

group, 283, 299 

hydroxide, 38, 110, 285 

hypochlorite, 110 

iodide, 285 

manganate, 369 

nitrate, 134, 286 

perchlorate, 87, 372 

permanganate, 43, 371 

sulphate, 289 
Propane, 385 
Propylene, 386 
Puddling, 361 
" Purple of Cassius," 381 
Pyrite, 367 
Pyroxylin, 406 

Quartation, 380 
Quartz, 258, 260 
Quartzite, 258, 260 
Quinine, 412 

Radicals, 400 
Reduction, 62, 167, 186 
Residues, 400 
Respiration, 169, 173 
Rhodium, 376 
Rock crystal, 260 
Rubidium, 300 
Ruby, 342 
copper, 325 



434 



INDEX, 



Ruthenium, 376 

Safety-lamp, 183, 387 
Sal ammoniac, 134, 296 
Salt, common, 97, 104, 290, 293, 
294 

Epsom, 321 

Glauber's, 292 

microcosmic, 299 
Saltpetre, 134, 286 

Chili, 134, 143, 291 
Salts, 118, 120, 122, 241, 271 

acid, 241 

decomposition of, 271 

ethereal, 397 

neutral, 241 

nomenclature of, 122 

normal, 241 
Saponification, 398 
Sapphire, 342 
Scandium, 214, 348 
Selenium, 241 
Serpentine, 319 
Siderite, 358 
Silica, 258, 260 
Silicates, 281 
Silicic acid, 259 

anhydride, 260 
Silicon, 258 

dioxide, 260 

fluoride, 224, 259 

hydride, 259, 260 
Silver, 332 

alloys, 334 

compounds, 334, 336, 417 

chloride, 334, 336 

nitrate, 334 
Slow oxidation, 33 
Smithsonite, 321 
Soaps, 394 
Soapstone, 319 
Soda, caustic, 38, 130. 291 

manufacture, 293, 294 
" Soda-water," 172 
Sodium, 37, 89, 289 

acid carbonate, 295 

ammonium phosphate, 299 

borate, 256, 296 

carbonate, 292 

chloride, 97, 104, 290, 293 

chlorplatinate, 378 



! Sodium compounds, 290, 424 

hydroxide, 38, 130, 291 

" hyposulphite, " 292 

metaphosphate, 281 

nitrate, 134, 143, 291 

phosphate, 281, 295 

pyrophosphate, 281 

sulphate, 97, 107, 292, 293 

tetraborate, 256, 296 

tbiosulphate, 292 
"Soft-soap," 396 
Solder, soft, 354 
Solution, 65, 66 
Solvay process, 294 
Sorrels, 396 
Spathic iron, 358 
Specific heat of metals, 336 
Spectroscope, 301 
Spelter, 322 
Sphalerite, 321 
Spiegel-iron, 360 
Spinel, 343 
" Spirits of hartshorn," 136 

of wine, 389 
Spiritus fumans Libavii, 356 
Stan nates, 355 
Stannic acid, 355 

chloride, 356 

hydroxide, 355 

oxide, 355 

sulphide, 356 
Stannous chloride, 355 

compounds, 354 
Starch, 382, 402, 406 
Stearin, 394 
Steel, 361 
Stibine, 252 
Strass, 315 
Strontium, 317 

Substituting power of the ele- 
ments, 204 
Substitution, 89, 387 
"Sugar of lead," 352, 393 
Sugar-refining, 404 
Sugars, 402 
Sulphates, 276 
Sulphides, 226, 230, 233, 273 
Sulphites, 279 
Sulphur, 26, 88, 226 

dimorphism of, 229 

dioxide, 235 



INDEX. 



435 



Sulphur group, 226 

trioxide, 235 
Symbols, 18, 19 
Synthesis, 49 

of water, 51, 52, 53 

Tannin, 410 

Tanning, 410 
Tantalum, 254 
Tartar, cream of, 397 

emetic, 252 
Tellurium, 241, 378 
Thallium, 339 

Thomas-Gilchrist process, 362 
Thorium, 260 
Tin, 15, 146, 353 

alloys, 354 

amalgam, 354 

compounds, 354, 357, 417 
Titanium, 260 
Toluene, 383, 386, 408 
Tungsten, 19 
Turpentine, 100 

Ultramarine, 345 
Uranates, 375 
Uranium, 375 
Uranyl, 375 
compounds, 375 

Valence, 203 
Vanadium, 254 
Vapor-densities, 195 
Verdigris, 393 
Vitriol, blue, 329 
green, 366 



Vitriol, oil of, 239 
white, 323 

Water, 47, 65 
analysis of, 50 
of crystallization, 47 
hard, 177, 309, 311, 395 
maximum density of, 65 
solvent properties, 65 
synthesis of, 51, 52, 53 
uses of, in the laboratory, 66 

11 Water-gas," 40, 91, 179 

Waters, natural, 63 

Water-vapor in the air, 129, 131 

Weldon's process, 370 

Wood-spirit, 383, 389 

Wood-vinegar, 393 

Wrought-iron, 361 

Xylene, 383, 386 

Yeast, 389 
Ytterbium, 339 
Yttrium, 339 

Zinc, 41, 42, 321 

alloys, 323 

chloride, 42, 109, 324 

compounds, 324, 421 

dust, 322 
«f method, 333 

oxide, 109, 323 

sulphate, 42, 323 
Zinc-white, 323 
Zircon, 260 
Zirconium, 260 



30 









